AGRIC.  DEPT, 


QUALITATIV  CHEMICAL  ANALYSIS 


A.  A.  NOYES 


A  COURSE  OF  INSTRUCTION 

IN  THE 

f 

QUALITATIV 
CHEMICAL   ANALYSIS 

OF  INORGANIC  SUBSTANCES 


BY 


ARTHUR  A.  NOYES 

PROFESSOR    OF    THEORETICAL    CHEMISTRY    IN    THE 
MASSACHUSETTS    INSTITUTE    OF    TECHNOLOGY 


SIXTH  EDITION 


NEW  YORK: 

THE  MACMILLAN  COMPANY. 

LONDON:  MACMILLAN  &  Co.,  LTD. 

1914 


CJ 


COPTBIGHT,  1914 

BY  ARTHTJK  A.  NOTES 


PREFACE. 


This  text-book  is  an  attempt,  on  the  experimental  side,  to  train  the  student 
of  qualitativ  analysis  in  careful  manipulation  and  exact  methods  of  procedure, 
such  as  are  commonly  employed  in  quantitativ  analysis.  It  is  an  attempt,  on 
the  theoretical  side,  to  make  clear  to  the  student  the  reason  for  each  operation 
and  result,  and  to  accustom  him  to  apply  to  them  the  laws  of  chemical  equilib- 
rium and  the  principles  relating  to  the  ionization  and  complex-formation  of 
substances  in  solution.  It  is  believed  that  in  both  these  ways  the  educational 
value  of  the  subject  is  greatly  increased. 

The  book  is  divided  into  two  main  Parts,  entitled  The  Course  of  Instruction 
and  The  System  of  Analysis.  In  presenting  the  System  of  Analysis  the  plan 
adopted  in  the  earlier  editions  of  this  book  has  been  followed,  namely,  that  of 
separating  sharply  the  description  of  the  operations  from  the  discussion  and 
explanation  of  them.  The  operations  are  described  with  as  great  definitness 
as  possible  in  short  paragraphs  entitled  "Procedures";  and  each  of  these  is 
followed  by  "Notes"  in  which  are  given  the  reasons  for  the  operations,  the 
precautions  necessary  and  difficulties  encountered  in  special  cases,  the  chemical 
behavior  of  the  different  elements,  the  indications  afforded  of  their  presence, 
and  the  application  of  the  theoretical  principles  to  the  reactions  involved. 
The  system  of  procedure  has  been  thoroly  revised  as  a  result  of  the  extended 
investigations  made  in  the  laboratory  of  this  Institute  during  the  past  six  years 
and  described  in  volumes  29,  30,  31,  and  34  of  the  Journal  of  the  American 
Chemical  Society.  As  a  result  of  these  investigations,  in  which  the  author 
has  had  the  able  cooperation  of  Professor  W.  C.  Bray  and  Professor  E.  B.  Spear, 
the  process  of  analysis  has  been  made  much  more  reliable,  so  that  now  it  is 
possible  with  careful  manipulation  to  detect  one  milligram  of  any  constituent 
in  the  presence  of  500  milligrams  of  any  other  (except  in  a  few  combinations 
where  the  limit  of  detectability  is  two  milligrams) .  At  the  same  time  the  process 
has  on  the  whole  been  considerably  simplified.  The  larger  size  of  the  present 
edition  is  due,  not  to  greater  complexity  of  the  process,  but  to  the  inclusion 
in  the  Procedures  of  the  more  explicit  directions  necessary  to  secure  accuracy 
in  the  separations  and  reactions,  to  the  insertion  of  confirmatory  tests  for  most 
of  the  elements,  to  the  development  of  a  more  systematic  process  for  the  detec- 
tion of  the  acidic  constituents,  and  to  the  elaboration  of  the  notes  and  especially 
the  inclusion  in  them  of  the  theoretical  explanations  made  possible  by  the 
recent  development  of  our  knowledge  of  solutions. 


Vl  PREFACE 

The  Course  of  Instruction  includes  two  sections — one  entitled  Laboratory 
Experiments,  giving  the  directions  for  the  laboratory  work;  and  the  other 
entitled  Questions  on  the  Experiments,  consisting  of  a  series  of  questions  to  be 
studied  in  connection  with  the  class-room  exercises. 

The  laboratory  work  described  in  the  section  on  Laboratory  Experiments  is 
from  beginning  to  end  closely  correlated  with  the  systematic  scheme  of  analysis. 
For  experience  has  convinced  the  author  that  the  plan  followed  in  many  text- 
books of  requiring  the  student  to  study  the  separate  reactions  characteristic  of  the 
various  elements  before  undertaking  their  systematic  separation  is  highly  unsat- 
1sfactory.  However  valuable  the  knowledge  of  the  additional  reactions  may  be, 
it  is  found  in  practis  that  the  performance  of  such  a  large  number  of  independent* 
disconnected  experiments  makes  little  impression  on  the  student's  mind  and 
fails  to  awaken  his  interest  in  the  subject.  Qualitativ  analysis  affords  an  effectiv 
means  of  teaching  a  part  of  inorganic  chemistry  chiefly  because  it  unites  into  a 
connected  whole  a  great  variety  of  isolated  facts,  and  because  the  student  sees 
a  practical  use  of  the  information  presented  to  him;  but  these  advantages 
evidently  do  not  apply  to  facts  not  directly  related  to  the  process  of  analysis. 

The  Questions  on  the  Experiments  are  intended  to  aid  the  student  in  under- 
standing the  work  that  he  is  doing  in  the  laboratory  and  to  make  sure  that  he 
derives  from  the  subject  the  mental  traning  which  it  ought  to  afford.  The 
questions  are  in  large  part  of  such  a  character  that,  in  order  to  answer  them 
properly,  the  student  not  only  must  study  the  Notes  on  the  Procedures,  but 
also  must  giv  to  the  questions  some  independent  thought.  It  is  assumed  in 
these  questions,  as  well  as  in  the  Notes  on  the  Procedures,  that  the  student 
has  acquired,  in  his  previous  course  on  Inorganic  Chemistry,  a  general  knowledge 
of  the  mass-action  law  and  of  the  chemical  aspects  of  the  ionic  theory. 

The  best  plan  of  conducting  the  course,  when  circumstances  permit,  is  in 
the  author's  opinion  as  follows:  The  class,  if  large,  is  divided  into  sections 
of  from  15  to  25  students.  At  the  beginning  of  each  laboratory  exercise,  which 
should,  if  possible,  be  three  hours  long,  or  at  the  beginning  of  every  second  lab- 
oratory exercise,  the  instructor  holds  a  class-room  conference  with  each  section, 
at  which  the  experiments  to  be  next  made  are  discussed  in  outline,  and  those 
made  at  the  previous  exercises  are  reviewed  in  detail.  The  conferences  are 
carried  on  mainly  by  questioning  the  individual  students  and  by  encouraging 
them  to  ask  questions  as  to  matters  which  they  do  not  understand.  As  prepa- 
ration for  each  conference  the  student  is  expected  to  study  the  Table  in  which 
is  presented  the  outline  of  the  analysis  of  the  group  to  be  next  taken  up,  to 
study  the  Notes  on  the  Procedures  which  he  has  worked  through  just  pre- 
viously hi  the  laboratory,  and  (unless  the  course  is  an  elementary  one)  to  answer 
the  questions  on  the  corresponding  Experiments.  To  ensure  that  the  student 
givs  proper  study  to  the  subject,  a  short  written  exercise  may  well  be  held  at 
the  beginning  of  each  conference.  In  the  laboratory  work  the  class  may  be 
kept  nearly  together  by  giving  to  the  faster  working  students  additional 
unknown  solutions  on  each  group,  and  by  allowing  those  who  are  falling  behind 


PREFACE  vii 

to  omit  some  of  the  less  important  Experiments.  In  the  laboratory  great  stress 
is  laid  on  careful  work,  such  as  will  enable  the  proportions  of  the  various  con- 
stituents present  in  unknown  solutions  to  be  estimated  and  small  quantities  of 
them  to  be  detected.  An  effectiv  means  of  teaching  the  details  of  manipula- 
tion is  for  the  instructor  to  carry  through  in  the  lecture-room,  after  the 
students  have  had  a  little  experience  of  their  own  in  the  laboratory,  the  com- 
plete process  for  the  analysis  of  the  copper-group. 

Even  when  the  time  available  for  the  subject  of  qualitativ  analysis  does  not 
permit  of  so  complete  a  course  as  that  here  presented,  the  student  gets,  in  the 
author's  opinion,  a  better  training  by  working  through  selected  parts  of  an 
exact  scheme  of  analysis  carefully  and  thoroly  than  he  does  by  covering  the 
whole  of  an  elementary  scheme  superficially.  Experiments  that  may  be  well 
omitted  in  briefer  courses  are  indicated  by  asterisks  prefixed  to  the  description 
of  them  in  the  section  entitled  Laboratory  Experiments. 


PREFACE    TO    THE    FIFTH   EDITION. 

In  this  edition  important  changes  and  additions  have  been  made  as  follows. 
The  Procedures  for  the  detection  of  the  basic  constituents  and  the  Notes  upon 
them  have  been  improved  in  many  matters  of  detail;  and  there  have  been  intro- 
duced a  simpler  process  for  the  detection  of  nickel  and  cobalt  and  a  more  satis- 
factory one  for  the  analysis  of  the  alkali-group.  The  part  of  the  book  relating 
to  the  detection  of  the  acidic  constituents  has  been  entirely  rewritten;  and  a  more 
complete  and  more  instructiv  system  of  analysis  for  those  constituents  has  been 
presented.  Many  additional  tabular  outlines  have  been  included  to  assist  the 
beginner  in  grasping  the  general  plan  of  the  separations;  and  the  Questions 
on  the  Experiments  have  been  revised. 

In  making  this  revision  the  author  has  had  the  benefit  of  many  valuable  sug- 
gestions from  Professors  Henry  Fay,  W.  T.  Hall,  and  A.  A.  Blanchard  of  this 
Institute,  and  from  Prof.  G.  S.  Forbes  of  Harvard  University. 

ARTHUR  A.  NOTES. 

MASSACHUSETTS  INSTITUTE  OF  TECHNOLOGY, 
BOSTON,  August,  1914. 


TABLE  OF  CONTENTS. 


PART  I.    THE  COURSE  OF  INSTRUCTION.  PAGE 

LABORATORY  EXPERIMENTS 1 

QUESTIONS  ON  THE  EXPERIMENTS 1$ 

PART  H.    THE  SYSTEM  OF  ANALYSIS. 

PREPARATION  OF  THE  SOLUTION 27 

DETECTION  OF  THE  BASIC  CONSTITUENTS 41 

DETECTION  OF  THE  ACIDIC  CONSTITUENTS 92 

APPENDICES. 

PREPARATION  OF  THE  REAGENTS 117 

PREPARATION  OF  THE  TEST-SOLUTIONS 120 

APPARATUS  REQUIRED 122 

IONIZATION-VALUES 123 

ATOMIC  WEIGHTS  OF  THE  COMMON  ELEMENTS 124 

SOLUBILITIES  OF  SLIGHTLY  SOLUBLE  SUBSTANCES 124 


PART  I. 
THE   COURSE  OF  INSTRUCTION, 

LABORATORY  EXPERIMENTS. 


DETECTION   OF   THE    BASIC   CONSTITUENTS. 

Preliminary  Work. — Check  off  on  an  apparatus  slip  (corresponding 
to  that  printed  on  page  122)  the  apparatus  found  in  the  desk,  and  sign 
and  hand  in  the  slip. 

Make  a  750  cc.  wash-bottle,  taking  pains  to  bend  the  tubes  and  to 
cut  them  off  so  as  to  correspond  closely  with  the  model  exhibited  in 
the  laboratory.  Make  also  a  250  cc.  wash-bottle  (for  washing  with 
hot  water  and  special  solutions). 

Make  a  dropper  about  10  cm.  (4  inches)  long  by  drawing  out  one 
end  of  a  glass  tube  to  a  fairly  wide  capillary  and  slightly  expanding 
the  other  end  with  the  aid  of  a  file  while  it  is  heated  in  a  flame.  Cap 
the  expanded  end  with  a  rubber  nipple. 

Make  3  stirring-rods  about  15  cm.  long  by  cutting  a  piece  of  glass 
rod  into  sections  and  then  rounding  the  ends  in  a  flame. 

Experiment  i. — Separation  of  the  Basic  Constituents  into  Groups. — 
In  connection  with  this  experiment  read  the  General  Discussion  on 
page  41  and  refer  to  Table  II  (page  42).  Measure  out  with  the  aid 
of  a  10  cc.  graduate  5  cc.  portions  of  the  test-solutions  (see  Note  1) 
of  AgN03,  Cu(N03)2,  Zn(N03)2,  Ca(N03)2,  and  KN03.  Mix  the  por- 
tions in  a  conical  flask,  add  5  cc.  6-normal  HN03  and  10  cc.  NH4C1 
solution,  shake  for  a  minute  or  two,  and  filter.  Dilute  the  filtrate 
with  water  to  a  volume  of  100  cc.  Place  it  in  a  200  cc.  conical  flask; 
insert  a  two-hole  rubber  stopper  through  which  passes  a  tube  leading 
to  the  bottom;  and  pass  in  a  slow  current  of  H2S,  until,  upon  shutting 
off  the  gas  and  shaking  thoroly,  the  liquid  smells  strongly  of  it.  Fil- 
ter. To  the  filtrate  add  10  cc.  NH4OH  and  3-5  cc.  (NH4)2S.  Shake 
the  mixture  and  filter.  Evaporate  the  filtrate  to  a  volume  of  about 
10  cc.,  filter,  and  to  the  cold  solution  add  15  cc.  (NH4)2CO3  reagent 
and  15  cc.  alcohol. 

In  this  experiment  and  all  subsequent  ones  observe  carefully  every- 
thing that  happens,  and  record  it  clearly  and  neatly  in  the  note-book 
in  the  way  described  in  Note  2,  together  with  the  equations  expressing 
all  the  chemical  changes  that  take  place.  In  connection  with  each  Ex- 
periment study  the  Table  and  the  Notes  referred  to  in  the  directions. 

1 


'    LABORATORY  EXPERIMENTS. 

Notes. — 1.  The  solutions  of  constituents  to  be  tested  for,  here  called  the 
test-solutions,  are  all  so  made  up  as  to  contain  10  mg.  (10  milligrams)  of  the 
constituent  per  cubic  centimeter  of  solution.  The  mixture  used  hi  this  ex- 
periment therefore  contains  50  mg.  of  each  of  the  basic  constituents  silver, 
copper,  zinc,  calcium,  and  potassium.  The  student  should  acquire  the 
habit  of  working  with  definit  quantities  of  the  constituents  and  of  noting  the 
size  of  the  precipitates  which  they  yield.  For  a  good  qualitativ  analysis  should 
not  only  show  the  presence  or  absence  of  the  various  constituents,  but  should 
also  furnish  an  estimate  of  the  proportions  in  which  they  are  present. 

Test-solutions  should  not  be  used  in  place  of  reagents,  nor  reagents  in  place  of 
test-solutions,  since  the  concentrations  are  as  a  rule  quite  different.  In  regard 
to  the  concentrations  of  reagents,  see  Note  3,  page  44. 

2.  In  the  note-book  the  operations  should  be  indicated  very  briefly;  but 
ererything  that  happens  should  be  recorded  fully,  tho  concisely.  Thus  the 
report  of  the  first  experiment  may  be  made  in  the  following  form: 

Expt  i. — Added  HN03:  no  change  observed. 
Added  NH4C1:  white  curdy  ppt. 

Ag+NO3-+NH4+Cl- = AgCH-NH4+NO3-. 
Passed  in  H$:  large  black  flocculent  ppt. 
Cu++(NO3-)2+H2S  =  CuS+2H+N03-. 


Solid  substances  involved  in  chemical  reactions  should  be  indicated  by  under- 
lining their  formulas.  Largely  ionized  dissolved  substances  should  be  written 
with  +  and  —  signs  attached  to  the  formulas  in  such  a  way  as  to  show  the 
ions  into  which  they  dissociate.  Slightly  ionized  dissolved  substances  should 
be  distinguished  by  not  attaching  these  signs  to  the  formulas.  In  regard  to 
the  ionization  of  substances,  see  the  Table  on  page  123. 

The  neatness,  accuracy,  and  completeness  of  the  notebook  record  will  be  an 
important  factor  in  determining  the  grade  of  the  student. 

Experiment  2. — Analysis  of  the  Silver-Group. — Mix  in  a  conical  flask 
20  cc.  of  the  test-solution  of  Pb(NOa)2  with  5  cc.  portions  of  the  test- 
solutions  of  Bi(NO3)3,  AgNO3,  and  Hg2(N03)2,  and  treat  the  mixture 
by  P.  11-15  (i.e.,  by  Procedures  11-15  of  the  System  of  Analysis  de- 
scribed in  Part  II,  on  pages  43-46);  omit  the  confirmatory  tests 
mentioned  at  the  end  of  P.  12  and  13.  Study  Table  III  (page  43) 
before  carrying  out  this  experiment;  and  in  connection  with  it  read  the 
Notes  on  P.  11-15. 

Note. — In  all  these  "Experiments"  only  the  Procedures  actually  named  in 
the  directions  should  be  worked  through,  omitting  any  others  that  may  be  inci- 
dentally referred  to  in  the  System  of  Analysis. 

Experiment  3. — Precipitation  by  Hydrogen  Sulfide. — To  10  cc. 
of  the  test-solution  of  Bi(N03)3  in  a  200  cc.  conical  flask  add  5  cc. 
HNO3,  10  cc.  NH4C1  solution,  and  75  cc.  water;  and,  without  filtering 
the  mixture,  pass  H2S  into  it  till  it  becomes  saturated,  in  the  way 


LABORATORY  EXPERIMENTS.  3 

described  in  the  first  paragraph  of  p.  21,  omitting  the  filtration  at  the 
end.  (The  HNO3  and  NH4C1  are  added  and  the  solution  is  diluted  to 
100  cc.  so  as  to  have  the  same  conditions  as  in  an  actual  analysis.)— 
Read  Notes  1  and  2,  P.  21. 

*To  10  cc.  of  the  test-solution  of  H3As04  add  5  cc.  HNO3,  10  cc. 
NH4C1  solution,  and  75  cc.  water.  Treat  this  solution  by  the  whole 
of  P.  21.— Read  Note  3,  P.  21. 

*  Note. — Experiments  or  parts  of  experiments  preceded  by  an  asterisk  may 
be  omitted  in  brief  courses  on  the  subject  when  the  instructor  so  directs. 

Experiment  4. — Effect  of  Acid  on  the  Precipitation  by  Hydrogen 
Sulfide. — Introduce  into  each  of  three  test-tubes  by  means  of  a  dropper 
(see  Note  2,  P.  11)  3  drops  of  the  test-solution  of  Cd(N03)2.  Add 
to  the  first  tube  1  cc.  HC1,  to  the  second  3  cc.  HC1,  and  to  the  third 

9  cc.  HC1.     Then  add  to  each  solution  enough  water  to  make  the 
volume  about  20  cc.,  and  pass  a  slow  current  of  H2S  into  it  for  about 
a   minute. — Repeat   the   last  test    (with   9    cc.    HC1),   substituting 
Cu(NO3)2  for  the  Cd(N03)2. — Calculate  the  normal  concentration  of 
the  HC1  in  each  tube  and  the  number  of  milligrams  of  cadmium  or 
copper  per  100  cc.  solution;  and  record  the  values  in  the  note-book. 
—Read  Notes  4  and  5,  P.  21. 

Experiment  5. — Effect  of  Oxidizing  Substances  on  Hydrogen  Sulfide. 
—To  20  cc.  of  the  test-solution  of  Fe(N03)3  add  10  cc.  NH4C1  solution; 
5  cc.  HN03,  and  65  cc.  water,  and  pass  in  H2S  till  the  solution  is  satu- 
rated. Repeat  this  experiment,  substituting  20  cc.  of  the  test-solution 
of  K2Cr04  for  that  of  the  Fe(N03)3 .— Read  Notes  6  and  7,  P.  21. 

Experiment  6. — Analysis  of  the  Copper-Group. — Mix  10  cc.  portions 
of  the  test-solutions  of  Hg(NO3)2,  Pb(NO3)2,  Bi(NO3)3,  Cu(NO3)2, 
and  Cd(N03)2,  add  5  cc.  HN03,  10  cc.  NH4C1  solution,  and  35  cc. 
water,  treat  the  mixture  by  P.  21,  and  treat  the  precipitate  so  obtained 
by  P.  31-38. — Refer  to  Table  V  (page  54),  and  read  the  Notes  on  P. 
31-38. 

Experiment  7. — Analysis  of  an  Unknown  Solution  for  Elements  of 
the  Copper-Group. — Ask  the  instructor  for  an  unknown  solution  con- 
taining elements  of  the  copper-group  ("unknown  A"),  and  analyze 

10  cc.  of  it  by  P.  21  and  P.  31-38,  first  adding  5  cc.  HN03  and  10  cc. 
NH4C1  solution.     Record  and  report  the  results  as  described  in  the 
following  Note.     Keep  all  final  tests  in  properly  labelled  test-tubes  or 
flasks  until  the  report  on  the  analysis  has  been  returned  by  the  in- 
structor. 


LABORATORY  EXPERIMENTS. 

Note. — Record  the  results  of  the  analyses  of  unknown  solutions  in  the  note- 
book in  three  columns  headed  respectivly,  "Operations,"  "Observations," 
"Conclusions."  The  operations  and  observations  are  to  be  recorded  in  the  same 
brief  form  employed  in  the  experiments  with  known  solutions.  In  the  column 
headed  Conclusions  are  to  be  inserted  the  conclusions  that  may  be  drawn  from 
each  observation  as  to  the  presence  or  absence  of  any  of  the  constituents  that 
may  be  present  in  the  unknown  solution.  The  chemical  equations  involved 
need  not  be  written.  Sum  up  at  the  end  the  constituents  that  have  been  found 
to  be  present,  giving  also  a  rough  estimate  of  the  quantity  of  each  of  them  per 
10  cc.  of  solution.  Quantities  less  than  5  mg.  may  be  reported  as  "small"; 
those  from  5  to  50  mg.  as  "medium";  and  those  greater  than  50  mg.  as  "large." 
(It  is  to  be  noted,  since  one  gram  of  a  non-metallic  solid  substance  is  ordinarily 
taken  for  analysis,  that  5  mg.  corresponds  to  the  presence  of  0.5%  and  50  mg.  to 
the  presence  of  5%  of  the  constituent  in  such  a  substance.)  The  quantity  of 
any  constituent  present  is  to  be  estimated  from  the  size  of  the  precipitate 
obtained  in  the  confirmatory  test  or  in  the  Procedure  preceding  it.  The  student 
should  make  it  a  habit  to  compare  this  precipitate  in  any  doubtful  case  with  that 
obtained  by  subjecting  a  known  quantity  of  the  test-solution  to  the  same  final 
Procedure.  For  this  purpose  the  test-solution  may  be  measured  out  with  the 
aid  of  a  dropper,  noting  that  three  medium-size  drops  correspond  to  about  0.1  cc. 
of  the  solution  or  to  1  mg.  of  the  constituent  contained  in  it. — The  instructor  will 
return  the  report  of  the  student  with  an  entry  upon  it  showing  the  quantities 
of  the  various  constituents  which  the  unknown  actually  contained. 

The  correctness  of  the  results  obtained  in  the  analysis  of  these  unknown 
solutions  is  an  important  factor  in  determining  the  grade  of  the  student. 

Experiment  8. — Behavior  of  Elements  of  the  Tin-Group  towards 
Hydrogen  Sulfide  and  Ammonium  Sulfide. — To  5  cc.  water  in  each  of 
three  test-tubes  add  from  a  dropper  6  drops  respectivly  of  the  test- 
solutions  of  AsCls,  of  SbCls,  and  of  SnCU.  (Note  that  6  drops  of  a 
test-solution  contains  2  mg.  of  the  constituent  to  be  tested  for.)  Pass 
H^S  into  each  tube  for  half  a  minute.  Then  add  from  a  graduate  2  cc. 
ammonium  polysulfide  solution.  Finally  add  3  cc.  HC1  slowly  to 
each  tube,  and  shake  the  mixture. — Compare  these  precipitates  with 
that  produced  by  mixing  5  cc.  water,  2  cc.  ammonium  polysulfide 
solution,  and  3  cc.  HC1,  and  shaking.  (In  an  actual  analysis  the  ana- 
lyst decides  from  the  appearance  of  the  HC1  precipitate  whether  it 
contains  an  appreciable  quantity  of  the  tin-group.) — Refer  to  Table 
IV  (page  47). 

Experiment  9. — Separation  of  the  Tin-Group  from  the  Copper-Group. 
— To  a  mixture  of  5  cc.  portions  of  the  test-solutions  of  Bi(NOa)s, 
AsCl3,  SbCl3,  and  SnCU  add  5  cc.  HNO3,  10  cc.  NH4C1  solution,  and 
enough  water  to  make  the  volume  100  cc.  Treat  the  mixture  by  the 
first  paragraph  of  P.  21,  filter  with  the  aid  of  suction  (see  Note  1,  P. 
23),  treat  the  precipitate  by  P.  22,  using  10  cc.  ammonium  polysulfide, 
reject  the  residue  of  Bi2S3,  and  treat  the  solution  by  P.  23.  Treat 


LABORATORY  EXPERIMENTS.  5 

at  once  the  precipitate  obtained  in  P.  23  as  described  in  Expt.  10. — 
Refer  to  Table  IV  (page  47);  and  read  the  Notes  on  P.  22  and  P.  23. 

Experiment  10. — Analysis  of  the  Tin-Group. — Treat  the  precipi- 
tated sulfides  obtained  in  Expt.  9  by  P.  41-46. — Refer  to  Table  VI 
(page  60),  and  read  the  Notes  on  P.  41-46. 

Experiment  n. — Analysis  of  Unknown  Solutions  for  Elements  of 
the  Copper  and  Tin  Groups. — Ask  for  two  unknown  solutions  contain- 
ing elements  of  these  groups  ("unknowns  B  and  C"),  and  analyze 
10  cc.  of  each  of  them.  First,  in  order  to  secure  the  proper  acid  con- 
centration for  the  H2S  precipitation,  make  the  solution  exactly  neutral 
by  adding  to  it  NH4OH  drop  by  drop  till  it  no  longer  reddens  blue 
litmus  paper,  and  add  just  5  cc.  HN03  and  enough  water  to  make  the 
volume  100  cc.  Then  treat  the  mixture  by  P.  21-46.  (Unknown  B 
will  contain  only  elements  of  the  tin-group.) 

Experiment  12. — Precipitation  of  the  Aluminum  and  Iron  Groups 
and  Solution  of  the  Group-Precipitate. — Treat  a  mixture  of  10  cc. 
portions  of  the  test-solutions  of  Co(N03)2  and  of  Fe(N03)  by  P.  51 
and  by  the  first  five  lines  of  P.  52.— Refer  to  Table  VII  (page  65), 
and  read  Note  1,  P.  51.  and  Notes  1-2,  P.  52. 

Experiment  13. — Behavior  of  Elements  of  the  Aluminum  and  Iron 
Groups  towards  Ammonium  Hydroxide  and  Sulfide. — To  5  cc.  portions 
of  the  test-solutions  of  A1(NO3)3,  CrCl8,  Fe(N03)3,  FeS04,  Zn(N03)«, 
Mn(N03)2,  Ni(NO3)2,  and  Co(NO3)2,  in  separate  test-tubes  add  3  cc. 
NH4C1  solution  and  8-10  drops  of  NH4OH,  and  note  the  result.  Then 
add  2-3  cc.  more  NH4OH.  Finally  add  1-2  cc.  (NH4)2S  to  each  tube. 
Filter  out  the  NiS  precipitate,  and  boil  the  filtrate  for  2  or  3  minutes. 
Record  the  results  of  all  these  tests  in  a  single  table,  so  as  to  show 
what  effect  is  observed  and  what  compound  is  formed  in  the  case  of 
each  element  upon  the  addition  of  each  reagent. — Study  the  results, 
refer  to  Table  VII,  and  read  Notes  2-5  and  8-10,  P.  51.  ' 

Experiment  14. — Behavior  of  Elements  of  the  Aluminum  and  Iron 
Groups  towards  Sodium  Hydroxide  and  Peroxide. — To  separate  5  cc. 
portions  of  the  test-solutions  named  in  Expt.  13,  add  8-10  drops  of 
NaOH,  and  note  the  result.  Then  add  2-3  cc.  more,  and  again  note 
the  result.  Finally  to  each  of  the  mixtures  add  gradually  from  a  dry 
7-cm.  test-tube  0.2-0.3  cc.  Na2O2  powder,  and  heat  it  to  boiling.  Re- 
cord all  the  results  in  a  single  table  as  in  Expt.  13. — Study  the  results, 
refer  to  Table  VII,  and  read  Notes  3-7,  P.  52. 

*Experiment  15. — Precipitation  of  Alkaline- Earth  Elements  by 
Ammonium  Hydroxide  in  the  Presence  of  Phosphate. — Heat  about 


6  LABORATORY  EXPERIMENTS. 

0.3  g.  solid  Ca3(P04)8  with  10  cc.  water;  then  add  5  cc.  HN03.  To 
the  solution  add  NEUOH  till  the  mixture  after  shaking  smells  of  it; 
filter  out  the  precipitate;  and  add  1-2  cc.  (NH4)2C03  reagent  to  the 
filtrate.— Read  Notes  6-7,  P.  51,  and  Note  8,  P.  52. 

Experiment  16. — Analysis  of  the  Aluminum-Group. — Treat  a  mix- 
ture of  10  cc.  portions  of  the  test-solutions  of  A1(N03)3,  CrCla,  and 
Zn(N03)2  by  the  second  paragraph  of  P.  52  and  by  P.  53-57.  Refer 
to  Table  VIII  (page  71),  and  read  the  Notes  on  P.  53-57. 

Experiment  17. — Analysis  of  the  Iron-Group:  Separation  of  Man- 
ganese and  Iron. — Treat  a  mixture  of  10  cc.  portions  of  the  test- 
solutions  of  Mn(N03)2,  Fe(NO3)3,  Zn(N03)2,  Co(N03)2,  and  Ni(N03)2 
by  the  second  paragraph  of  P.  52;  and  treat  the  precipitate  there- 
by obtained  by  P.  61,  62,  64,  and  66.  Treat  the  precipitate  containing 
the  zinc,  cobalt,  and  nickel  as  described  in  Expt.  18. — Refer  to  Table 
IX  (page  74),  considering  only  the  case  where  " phosphate  is  absent," 
and  read  the  Notes  on  P.  61,  62,  64,  and  66. 

*Experiment  18. — Analysis  of  the  Iron-Group:  Separation  of  Zinc, 
Nickel,  and  Cobalt. — Treat  the  residue  and  precipitate  obtained  in 
Expt.  17  by  P.  67-70.— Refer  to  Table  X  (page  78),  and  read  the 
Notes  on  P.  67-70. 

Note. — In  brief  courses  this  experiment  may  be  omitted;  and  in  the  subse- 
quent analyses  of  unknowns  cobalt  and  nickel  may  be  detected  (without  dis- 
tinguishing them)  by  the  formation  of  a  black  precipitate  in  P.  66,  and  zinc 
may  be  tested  for  only  in  the  aluminum-group. 

*Experiment  19. — Modification  of  the  Analysis  of  the  Iron-Group  in 
the  Presence  of  Phosphate  for  the  Purpose  of  Detecting  Alkaline- Earth 
Elements. — Mix  together  10  cc.  portions  of  the  test-solutions  of 
Fe(N03)3,  of  Co(N03)2,  and  of  Caa(PO4)s  in  HN03.  Treat  one-tenth 
of  this  solution  by  P.  63,  and  treat  the  remainder  of  it  by  P.  65  and 
66.  To  the  filtrate  obtained  in  P.  66  add  2-3  cc.  (NH4)2C03  reagent. 
—Read  the  Notes  on  P.  65. 

Experiment  20. — Analysis  of  Unknown  Solutions  for  Elements  of 
the  Aluminum  and  Iron  Groups. — Ask  the  instructor  for  two  unknown 
solutions  for  this  purpose  ("unknowns  D  and  E"),  and  treat  10  cc. 
of  each  by  P.  51-57  and  61-70.  (Unknown  E  will  contain  phosphate.) 

Experiment  21. — Precipitation  of  the  Alkaline- Earth  Group. — To 
2  cc.  of  the  test-solution  of  Mg(N03)2  add  10  cc.  water  and  1-2  cc. 
(NH4)2C03  reagent,  and  shake  the  mixture  for  about  a  minute.  Then 
add  (in  accordance  with  P.  81)  15  cc.  (NH4)2C03  reagent  and  15  cc. 
95  per  cent  alcohol,  and  shake  for  a  minute  more. 


LABORATORY  EXPERIMENTS.  7 

To  2  cc.  of  the  test-solution  of  Ca(N03)2  add  10  cc.  water  and  2  cc. 
(NH4)2C03  reagent,  shake,  and  after  2-3  minutes  filter  out  the  pre- 
cipitate. To  the  filtrate  add  15  cc.  (NH^COs  reagent  and  15  cc. 
95  per  cent  alcohol. — Read  the  Notes  on  P.  81. 

Experiment  22. — Analysis  of  the  Alkaline- Earth  Group. — Mix  to- 
gether in  a  small  flask  3  cc.  portions  of  the  test-solutions  of  BaCl2, 
Sr(N03)2,  Ca(N03)2,  and  Mg(N03)2.  Add  to  the  mixture  30  cc.  of  the 
(NH4)2C03  reagent  and  30  cc.  95  per  cent  alcohol,  and  shake  it  for 
about  5  minutes.  Filter  out  the  precipitate  and  treat  it  by  P.  82-89. 
—Refer  to  Table  XI  (page  81;,  and  read  the  Notes  on  P.  82-89. 

Experiment  23. — Analysis  of  the  Alkali-Group. — Mix  together  10  cc. 
portions  of  the  test-solutions  of  KNOs  and  NaN03,  add  10  cc.  NH4C1 
solution  and  10  cc.  (NH4)2C03  reagent,  and  treat  the  mixture  by  P.  91- 
95.— Refer  to  Table  XII  (page  86),  and  read  the  Notes  on  P.  91-95. 

Experiment  24. — Analysis  of  an  Unknown  Solution  for  Elements  of 
the  Alkaline -Earth  and  Alkali  Groups. — Ask  the  instructor  for  an  un- 
known solution  for  this  purpose  ("unknown  F"),  and  analyze  10  cc. 
of  it  by  P.  81-89  and  P.  91-95. 

*Experhnent  25. — Detection  of  Ammonium. — Treat  0.2-0.3  g.  solid 
NH4C1  by  P.  96.— Read  the  Notes  on  P.  96. 

*Experiment  26. — Determination  of  the  State  of  Oxidation  of  Iron. 
—Treat  0.3  g.  finely  powdered  Fe3O4  by  P.  97,  omitting  tests  a,  6, 
and  c. — Read  the  Notes  on  P.  97. 

*Experiment  27. — Determination  of  the  State  of  Oxidation  of  Arsenic. 
—Treat  1  cc.  of  the  test-solution  of  H3As04  by  P.  98.— Read  the 
Notes  on  P.  98. 

*Expenment  28. — Detection  of  a  Small  Quantity  of  Arsenic. — Add 
6  drops  of  the  test-solution  of  H3As04  to  10  cc.  H2S04,  and  treat  the 
mixture  by  P.  99.— Read  the  Notes  on  P.  99. 

Experiment  29. — Analysis  of  Unknown  Solutions  for  All  the  Basic 
Constituents. — Ask  the  instructor  for  two  unknown  solutions  for  this 
purpose  ("unknowns  G  and  H")}  and  analyze  10  cc.  of  each  of  them 
by  P.  11-95.  Before  precipitating  with  H2S,  exactly  neutralize  the 
solution  with  NH4OH  and  add  5  cc.  HN03. 

Note. — In  complete  analyses  of  this  kind  where  a  number  of  different  precipi- 
tates and  filtrates  are  sue  essivly  obtained,  any  of  these  that  are  set  aside, 
even  temporarily,  should  be  distinctly  labelled,  in  order  to  avoid  mistakes. 
A  convenient  method  of  doing  this  is  to  mark  on  the  label  simply  the  Procedure 
by  which  the  precipitate  or  filtrate  is  next  to  be  treated;  thus  the  H2S  precipi- 
tate would  be  marked  P.  22,  and  the  filtrate  from  it  P.  51.  The  final  tests  for 
any  element  may  be  marked  Test  for  Pb,  Test  for  Al,  etc. 

2 


8  LABORATORY  EXPERIMENTS. 

DETECTION   OF  THE   ACIDIC   CONSTITUENTS. 

Experiment  30. — Detection  of  Readily  Volatil  Constituents. — Test  a 
mixture  of  5  drops  of  the  test-solution  of  Na2CO3  and  5  drops  of  the 
test-solution  of  Na2S  (in  place  of  "0.3  g.  of  the  finely  powdered  sub- 
stance") for  carbonate  and  sulfide  by  P.  101.  Test  5  drops  of  the 
test-solution  of  NaN02  for  nitrite  by  P.  101.  Test  5  drops  of  the 
test-solution  of  KCN  for  cyanide  by  P.  101. — Read  the  General 
Discussion  on  pages  93-94,  refer  to  Table  XIII  page  96),  and  read 
the  Notes  on  P.  101. 

Experiment  31. — Detection  of  Sulfate,  Sulfite,  and  Fluoride. — Mix 
together  2  cc.  portions  of  the  test-solutions  of  NajSCX,  Na^Os,  and 
KF,  add  2  cc.  HN03,  and  treat  the  mixture  by  P.  102.  (Reject  the 
CaCl2  precipitate.) — Refer  to  Table  XIV  (page  97),  and  read  the 
Notes  on  P.  102. 

Experiment  32. — Detection  of  Halides  and  of  Chlorate. — Mix  to- 
gether 2  cc.  portions  of  the  test-solutions  of  KC1  and  KC103,  add 
2  cc.  HN03,  and  treat  the  mixture  by  P.  103.— Refer  to  Table  XIV 
(page  97),  and  read  the  Notes  on  P.  103. 

Experiment  33. — Detection  of  Phosphate. — To  2  cc.  of  the  test-solu- 
tion of  Na2HPO4  add  2  cc.  HN03,  and  treat  the  mixture  by  P.  104. 
—Read  the  Notes  on  P.  104. 

Experiment  34. — Detection  of  Thiocyanate. — To  2  cc.  of  the  test- 
solution  of  KSCN  add  10  drops  of  HNOs,  and  treat  the  mixture  by 
P.  105.— Read  the  Notes  on  P.  105. 

Experiment  35. — Detection  of  the  Separate  Halides. — Mix  together 
2  cc.  portions  of  the  test -solutions  of  KC1,  KBr,  and  KI,  and  treat 
the  mixture  by  P.  106.— Refer  to  Table  XV  (page  100),  and  read 
the  Notes  on  P.  106. 

Experiment  36. — Analysis  of  an  Unknown  Solution  for  the  Acidic 
Constituents  tested  for  in  Nitric-Acid  Solution. — Ask  the  instructor 
for  an  unknown  solution  for  this  purpose  ("unknown  J").  To  10 
cc.  of  it  add  15  cc.  water  and  5  cc.  HNOs,  and  treat  portions  of  the 
mixture  by  P.  102-106.— Refer  to  Tables  XIV  and  XV  (pages  97 
and  100). 

*Experiment  37. — Distillation  with  Phosphoric  Acid  and  Detection  of 
Carbonate,  Halides,  and  Sulfate. — Mix  3  cc.  portions  of  the  test-solu- 
tions of  Na2C03,  KC1,  and  K2S04,  and  treat  this  mixture  (in  place  of 
"2  g.  of  the  finely  powdered  substance")  by  P.  111.  Treat  a  portion 
of  the  second  distillate  by  the  first  paragraph  of  P.  116.  Treat  the 


LABORATORY  EXPERIMENTS.  9 

third  distillate  by  P.  119.— Refer  to  Table  XVI  (page  102),  and  read 
the  Notes  on  P.  Ill,  116,  and  119. 

*Experiment  38. — Detection  of  Carbonate  and  Sulfite  in  the  Presence 
of  Each  Other. — Mix  3  cc.  portions  of  the  test-solutions  of  Na2C03  and 
Na2SOs,  add  to  the  mixture  10  cc.  Ba(OH)2  solution,  acidify  with 
HAc,  and  treat  the  mixture  (in  place  of  "one-half  of  the  first  dis- 
tillate") by  P.  112.  Refer  to  the  first  part  of  Table  XVII  (page  105), 
and  read  the  Notes  on  P.  112. 

*Experiment  39. — Detection  of  lodin-Liberating  Constituents. — Treat 
10  cc.  of  the  test-solution  of  NaOCl  (which  contains  also  an  equiva- 
lent quantity  of  NaCl  and  some  Na2C03)  by  the  first  paragraph  of 
P.  111.  Treat  the  whole  distillate  so  obtained  by  the  first  two 
paragraphs  of  P.  113. — Treat  3  cc.  of  the  test-solution  of  KNC>2  by 
the  first  two  paragraphs  of  P.  113. — Refer  to  the  middle  part  of 
Table  XVII  (page  105),  and  read  the  Notes  on  P.  113. 

*Experiment  40.— Detection  of  Cyanide. — Treat  3  cc.  of  the  test- 
solution  of  KCN  by  P.  115.— Refer  to  the  last  column  of  Table  XVII 
(page  105),  and  read  Notes  1-3,  P.  115. 

*Experiment  41. — Analysis  of  an  Unknown  Solution  for  Acidic  Con- 
stituents passing  into  the  Phosphoric  Acid  Distillates. — Ask  the  in- 
structor for  an  unknown  solution  for  this  purpose  ("unknown  K"), 
and  treat  10  cc.  of  it  by  P.  111.     Treat  the  distillates  by  P.  112-119. 
-Refer  to  Tables  XVII  and  XVIII  (pages  105  and  109). 

Experiment  42. — Detection  of  Borate. — Treat  0.2-0.3  g.  of  solid 
borax  (Na2B4O7)  by  P.  121. — In  connection  with  each  one  of  Expts. 
42-47  refer  to  Table  XIX  (page  111),  and  read  the  Notes  on  the  Pro- 
cedure involved  in  the  Experiment. 

Experiment  43. — Detection  of  Fluoride. — Treat  0.2-0.3  g.  of  solid 
CaF2  by  P.  122. 

Experiment  44. — Detection  of  Nitrate. — Treat  0.2-0.3  g.  of  solid 
KN03  by  P.  124. 

*Experiment  45. — Detection  of  Nitrite. — Treat  1  cc.  of  the  test- 
solution  of  KNO2  by  P.  125. 

*Experiment  46. — Detection  of  Hypochlorite. — Treat  3  cc.  of  the 
test-solution  of -NaOCl  by  P.  126. 

*Experiment  47. — Detection  of  Chlorate  in  Presence  of  Hypochlorite. 
— Mix  3  cc.  portions  of  the  test-solutions  of  NaOCl  and  KC103,  and 
treat  the  mixture  by  P.  127. 


10  LABORATORY  EXPERIMENTS. 

PREPARATION     OF    THE     SOLUTION    AND     COMPLETE   ANALYSES    OF 
UNKNOWN   SOLID   SUBSTANCES. 

Experiment  48. — Substances  Soluble  in  Water  or  Dilute  Acid. — Ask 
the  instructor  for  two  such  unknown  substances  ("unknowns  I  and 
II"),  and  treat  portions  of  each  of  them  by  P.  1,  by  P.  2  followed  by 
P.  11-95,  by  P.  96-98,  and  by  P.  100,  101-106,  121,  and  124-126.- 
Read  the  Notes  on  P.  1  and  2. — Record  and  report  the  results  in  the 
note-book  as  directed  in  the  Note  on  Expt.  7.  In  the  case  of  a  solid 
substance  not  only  report  the  constituents  and  the  proportions  of 
them  present,  but  state  the  compound  or  compounds  of  which  the 
substance  seems  to  be  mainly  composed. 

Note. — In  analyzing  unknown  solids  the  quantity  taken  for  the  analysis 
should  be  weighed  (within  0.1  g.)  on  a  rough  balance,  not  guessed  at  nor  esti- 
mated by  volume. 

Experiment  49. —  Non-Metallic  Substances  requiring  Treatment  with 
Concentrated  Acids. — Ask  the  instructor  for  two  such  substances 
("unknowns  III  and  IV"),  and  treat  portions  of  each  of  them  by  P.  1, 
by  P.  2-3  (or  2-4,  if  necessary)  followed  by  P.  11-95  (or  21-95),  by 
P.  96-98,  and  by  P.  100,  111-119,  and  121-127.  If  there  is  any 
undissolved  residue  (of  silica  or  silicate)  at  the  end  of  P.  4,  disregard 
it  in  these  analyses. — Refer  to  Table  I  (page  29),  and  read  Notes  1-4, 
P.  3,  and  Notes  1-2,  P.  4. 

Experiment  50. — Alloys  Dissolved  by  Concentrated  Acids. — Ask  the 
instructor  for  two  such  alloys  ("unknowns  V  and  VI") ,  and  treat  0.5  g. 
of  each  by  P.  3-4  and  P.  11-70  (or  P.  21-70).— Read  Notes  5-9,  P.  3, 
and  Notes  3-4,  P.  4. 

*Experiment  51. — Mineral  Substances  Not  Completely  Dissolved  by 
Concentrated  Nitric  and  Hydrochloric  Acids. — Ask  the  instructor  for 
two  such  substances  ("unknowns  VII  and  VIII").  Treat  1  g.  of 
each  of  them  by  P.  2-6,  followed  by  P.  11-95.— Refer  to  Table  I 
(page  29),  and  read  the  Notes  on  P.  5  and  6. — Treat  fresh  portions  of 
each  of  the  substances  by  P.  131. — Read  the  Notes  on  P.  131. 

Ask  the  instructor  for  another  such  substance  ("unknown  IX"). 
Treat  1  g.  of  it  by  P.  2-4.  Treat  the  solution  obtained  in  P.  4  by 
P.  21-95,  and  the  residue  obtained  in  P.  4  by  P.  7,  followed  by 
P.  21-89.  Reserve,  as  directed,  one-half  of  the  aqueous  extract  of  the 
fused  mass  obtained  in  P.  7,  and  test  it  for  acidic  constituents  as  de- 
scribed in  the  third  and  fourth  paragraphs  of  P.  131. — Read  the 
Notes  on  P.  7. — Treat  fresh  portions  of  the  substance  by  P.  101  and 
P.  102-104. 


LABORATORY  EXPERIMENTS.  11 

*Experiment  52. — Substances  Containing  Organic  Matter. — Ask  the 
instructor  for  such  a  substance  ("unknown  X")>  and  treat  portions  of 
it  by  P.  1,  by  P.  8  followed  by  P.  11-95,  by  P.  96-98,  and  by  P. 
101-106.— Read  the  Notes  on  P.  8. 


QUESTIONS  ON  THE  EXPERIMENTS. 


DETECTION    OF   THE   BASIC    CONSTITUENTS. 

Experiment  i. — 1.  In  precipitating  the  silver-group  in  an  actual  analysis  could 
the  NH4C1  be  replaced  by  NaCl?  by  HC1?  (In  the  case  of  all  questions  which  can 
be  answered  by  "yes"  or  "no,"  giv  the  reasons  for  the  answer.) 

2.  If  the  NH4C1  were  not  added,  what  would  happen  to  the  silver  in  the  subse- 
quent parts  of  the  experiment? 

3.  Of  the  five  basic  constituents  present  in  the  mixture  why  is  silver  the  only  one 
that  is  precipitated  by  NH4C1? 

4.  If  enough  H2S  were  not  used  to  precipitate  all  the  copper,  how  would  it  behave 
on  the  subsequent  addition  of  NH4OH  and  (NH4)2S? 

5.  What  is  the  first  reaction  that  takes  place  when  NH^OH  is  added  to  the  filtrate 
from  the  H2S  precipitate? 

6.  What  would  happen  to  the  (NH^S  if  it  were  added  directly  to  the  filtrate  from 
the  H2S  precipitate,  without  first  adding  NH4OH? 

7.  What  happens  to  the  (NH4)2S  when  the  filtrate  from  the  (NHOaS  precipitate 
is  evaporated? 

8.  If  all  the  basic  constituents  had  been  present  in  the  original  mixture  used  for 
this  experiment,  what  ones  would  have  been  precipitated  by  (a)  NH4C1,  (6)  H2S,  (c) 
NH4OH  and  (NH^S,  (d)  (NH4)2CO3?    (e)  What  ones  would  have  been  left  with  the 
potassium  hi  the  filtrate  from  the  (NH4)2COa  precipitate? 

Experiment  2. — 1.  What  would  be  meant  by  the  statement  that  a  certain  quan- 
tity of  lead  nitrate  is  equivalent  to  a  certain  other  quantity  of  ammonium  chloride? 

2.  In  making  up  one  liter  of  a  1-normal  solution  of  NH4C1,  how  many  grams  of 
the  salt  should  be  weighed  out,  and  how  much  water  should  be  added  to  it?    (For 
the  atomic-weight  values  needed  in  answering  this  and  other  questions  see  the  table 
on  page  124.) 

3.  In  making  up  a  liter  of  a  6-normal  solution  of  H2S04,  how  many  cubic  cen- 
timeters of  96%  sulfuric  acid  (s.  g.,  1.84)  should  be  used,  and  how  much  water 
should  be  added  to  it? 

4.  Approximately  how  many  cubic  centimeters  of  the  NH4C1  solution  would  be 
required  to  precipitate  500  mg.  of  silver?    (Calculate  first  the  number  of  equivalents 
corresponding  to  500  mg.  Ag. — Since  1  g.  of  the  unknown  substance  is  ordinarily 
taken  for  the  analysis  for  basic  constituents,  500  mg.  is  as  large  a  quantity  of  any 
element  as  is  likely  to  be  present.) 

5.  Why  is  a  considerable  excess  of  NH4C1  added?    (The  word  excess  signifies  the 
quantity  added  beyond  the  equivalent  quantity  theoretically  required  to  produce 
the  reaction  in  question.) 

6.  The  solubility  of  Pbd2  at  20°  in  water  is  0.070  equivalents  per  liter  and  in 
0.2-normal  NH4C1  solution  is  0.018  equivalents  per  liter.    Explain  by  the  solubility- 
product  principle  why  NH4C1  diminishes  the  solubility.    (This  question  may  be 
answered  by  shortening  the  complete  explanation  given  in  Note  6,  P.  11,  as  follows: 
"In  any  dilute  sol'n  satur.  with  PbCl2,   (Pb++)  X (Cl~)2  =  satur.  value.     NH4C1 
added  to  such  a  sol'n  causes,  owing  to  its  ionization  into  NH4+  and  Cl~,  an  increase 
in   (Cl~),  and  therefore  raises  (Pb++)  X  (Cl~)2  above  the  saturation  value,  BO  that 

13 


14  QUESTIONS  ON   THE  EXPERIMENTS. 

Pbdz  ppts."  (All  other  questions  as  to  the  effect  of  one  substance  on  the  solubility  of 
another  substance  should  be  answered  in  a  similar  way.  Always  consider  the  effect  which 
the  added  substance  may  have  on  the  concentration  of  each  of  the  ions  of  the  salt  with  which 
the  solution  is  saturated,  and  state  the  reason  for  any  such  effect.) 

7.  Calculate  by  the  solubility-product  principle,  from  the  fact  that  the  solubility 
of  PbCk  in  water  at  20°  is  0.070  normal,  what  its  solubility  would  be  in  a  solution 
0.16  normal  in  chloride-ion  (which  is  approximately  the  chloride-ion  concentration 
in  a  0.2  normal  NHjCl  solution).     Assume  the  PbCl2  to  be  completely  ionized. 

8.  From  the  data  given  in  Question  6  calculate  how  many  milligrams  of  lead 
would  have  to  be  present  in  40  cc.  water  at  20°,  in  order  that  any  precipitation  of 
PbCl2  may  result  on  adding  to  it  10  cc.  1— normal  NH^Cl  solution. 

9.  In  the  precipitation  of  bismuth  by  the  NKiCl  solution,  what  ion -concentration 
product  comes  into  consideration?    What  must  be  true  of  its  value  in  order  that 
bismuth  may  be  precipitated?    Explain  why  increasing  the  HNOa-concentration 
increases  the  quantity  of  bismuth  that  remains  in  solution. 

10.  The  solubility  of  AgCl  at  100°  is  0.022  g.  per  liter.     Calculate  how  many 
milligrams  of  silver  might  be  lost  if  the  chloride  precipitate  were  washed  with 
100  cc.  boiling  water. 

11.  In  testing  for  lead  in  P.  13,  explain  why  the  addition  of  an  excess  of  H2SO4  to 
the  solution  diminishes  the  solubility  of  the  precipitate,  and  thus  increases  the  deli- 
cacy of  the  test. 

12.  What  other  elements  besides  lead  might  be  precipitated  by  adding  H2S04  to 
a  solution  in  which  they  were  present?     Since  these  other  elements  are  precipitable 
by  H2SO4,  why  does  the  formation  of  a  precipitate  in  P.  13  show  the  presence  of 
lead? 

13.  Explain  by  the  solubility-product  principle  why  the  formation  of  the  complex 
salt  Ag(NHa)2^Cl~  causes  AgCl  to  be  much  more  soluble  in  NBUOH  solution  than 
hi  water.     (Answer  in  accordance  with  the  Note  on  Question  6.) 

14.  Formulate  the  mass-action  expression  for  the  equilibrium  between  the  com- 
plex cation  Ag(NH3)2+  and  its  constituents.     Show  by  reference  to  this  expression 
and  the  solubility-product  principle  why  the  addition  of  HNOs  causes  AgCl  to  be 
precipitated  out  of  its  solution  in  NH^OH. 

Experiments  3  ancl  4. — 1.  In  precipitating  with  E^S  in  P.  21  what  is  the  reason  for 
adding  5  cc.  HNOs  and  diluting  the  solution  to  100  cc.?  Why  not  use  less  acid  and 
thus  avoid  all  risk  of  failing  to  precipitate  the  elements  of  the  copper  and  tin  groups? 

2.  In  passing  H2S  into  a  Cu(NOa)2  solution,  at  what  stage  in  the  process  does 
the  solution  after  shaking  begin  to  smell  of  the  gas? 

3.  What  principle  determine  how  the  quantity  of  £[28  dissolved  by  a  given 
quantity  of  water  varies  with  its  partial  pressure?    What  would  its  partial  pressure 
be  in  a  mixture  made  by  mixing  1  volume  of  H^S  with  4  volumes  of  air,  each  at  a 
pressure  of  one  atmosphere? 

4.  Why  would  a  larger  quantity  of  an  element  have  to  be  present  in  order  to  giv 
a  precipitate  if  in  P.  21  the  solution  were  treated  with  H^S  in  an  open  beaker,  in- 
stead of  in  the  closed  flask? 

5.  Giv  two  reasons  why  a  larger  quantity  of  an  element  would  have  to  be  present 
to  giv  a  precipitate  if  the  solution  were  saturated  with  H2S  at  80°,  instead  of  at  20°. 

6.  The  solubility  (in  equivalents  per  liter)  of  freshly  precipitated  ZnS  in  water  is 
about  100  times  as  great  as  that  of  CdS.     Calculate  by  the  principles  discussed 
in  Note  4,  P.  21,  the  ratio  of  the  hydrogen-ion  concentrations  at  which  the  precip- 


QUESTIONS  ON  THE  EXPERIMENTS.  15 

itation  of  cadmium  and  zinc  will    barely  take  place  when  the  concentration  of 
each  of  them  has  any  definit  value  (for  example,  0.0001  equivalents  per  liter). 

7.  The  pressure* volume  relations  of  perfect  gases  are  expressed  03^  the  equation 
p  v  I  T  =  82  N,  when  the  pressure  p  is  in  atmospheres,  the  volume  v  in  cubic  centi- 
meters, and  the  temperature  T  in  centigrade-degrees  on  the  absolute  scale,  and 
when  the  quantity  of  the  gas  is  N  gram-molecular-weights.     Calculate  the  number 
of  cubic  centimeters  of  H^S  at  25°  required  to  precipitate  500  mg.  of  copper. 

8.  By  what  reaction  is  the  HNOs  destroyed  when  the  arsenic  solution  to  which 
HC1  has  been  added  is  evaporated  to  dryness?     Could  HC1  be  destroyed  in  the  same 
way  by  evaporating  a  solution  of  chloride  with  HNOs? 

9.  If  the  HNOs  were  not  so  destroyed,  what  would  happen  when  the  H^S  is 
passed  into  the  hot,  strongly  acid  solution? 

Experiment  5. — 1.  What  substances  besides  ferric  salts  might  be  present  which 
would  cause  precipitation  of  sulfur  in  P.  21? 

2.  Write  the  equation  expressing  the  reaction  between  each  of  these  substances 
and  H2S,  balancing  the  equations  by  the  method  described  in  Note  7,  P.  21. 

3.  Write  by  the  same  method  the  equations  expressing  the  oxidation  of  H^S,  in 
one  case  to  sulfur  and  in  another  to  H2S04,  by  hot  fairly  concentrated  HNOs, 
assuming  that  the  HNOs  is  reduced  to  NO. 

Experiment  6. — 1.  Make  a  table  showing  briefly  in  the  first  column  the  chemical 
operations  involved  in  analyzing  a  solution  for  mercury  and  lead  (by  P.  21  and  P. 
31-34),  and  showing  in  a  second  and  in  a  third  column  the  behavior  of  these  two 
elements  in  each  operation. — "Behavior"  in  this  and  later  questions  means  both  the 
effect  observed  and  the  chemical  compound  produced.  Thus,  the  first  two  operations 
and  the  results  of  them  should  be  entered  as  follows: 

Operation  Behavior  of  Mercury  Behavior  of  Lead 

Satur.  with  H2S.  Black  ppt.  of  HgS.  Black  ppt.  of  PbS. 

Boil  with  2-n.  HNOs.  No  change.  Ppt.  diss.,  forming  colorless 

sol'nofPb++(N03-)2. 

2.  Make  a  similar  table  showing  the  operations  involved  in  analyzing  a  solution 
for  bismuth,  copper,  and  cadmium  (by  P.  21,  31,  33,  and  35-38),  and  showing  the 
behavior  of  these  elements. 

3.  Explain  by  the  solubility-product  principle  the  fact  that  CuS,  which  is  only 
slightly  soluble  in  hot  dilute  HC1,  dissolves  readily  in  hot  dilute  HNOa  of  the  same 
concentration. 

4.  Write  the  equation  expressing  the  dissolving  of  CuS  in  the  HNOs. 

5.  Why  is  a  black  residue  left  undissolved  by  HNOs  not  sufficient  evidence  of  the 
presence  of  mercury? 

6.  Suggest  a  reason  why  hi  the  confirmatory  test  for  mercury  the  addition  of  HC1 
with  the  SnCl2  tends  to  prevent  the  immediate  reduction  of  the  Hg&Ck  to  Hg. 

7.  Why  does  the  evaporation  with  H2SO4  convert  the  salts  present  into  sulfates? 
Could  sulfates  be  converted  into  nitrates  by  evaporating  with  a  large  excess  of 
HN03? 

8.  Explain  with  reference  to  the  solubility-product  principle  why  PbS04  is  much 
more  soluble  in  dilute  HNOa  than  in  water.     (£[2804  in  dilute  solution  is  dissociated 
almost  completely  into  H+  and  HSO4~;  but  the  latter  ion  is  only  to  a  moderate  ex- 
tent dissociated  into  H+  and  SC>4~). 

9.  What  effect,  as  compared  with  that  of  HNOa,  would  HC1  have  on  the  solubility 
of  PbS04?    What  effect  would  KNO3  have?    Giv  reasons.     (K2S04  in  dilute  solu- 


16  QUESTIONS  ON  THE  EXPERIMENTS. 

tion,  like  other  unibivalent  salts,  but  unlike  £[2804,  is  almost  completely  dissociated 
into  the  simple  ions,  K+  and  SO4=,  with  formation  of  only  a  small  proportion  of  the 
intermediate  ion,  KSCV.) 

10.  Explain  by  the  solubility-product  principle  why  the  fact  that  PbAc2  is  a  slightly 
ionized  substance  should  cause  PbSO4  to  dissolve  much  more  readily  in  NEUAc  solu- 
tion than  hi  water. 

11.  Would  you  expect  PbCrO4  also  to  be  more  soluble  in  NELiAc  solution  than 
in  water?    Why  or  why  not?   if  so,  why  does  PbCrO4  precipitate  from  the  same 
NH4Ac  solution  that  dissolves  PbSO4? 

12.  Arrange  all  the  compounds  of  lead  thus  far  met  with  in  the  order  in  which  their 
solubility  in  water  decreases. 

13.  If  in  an  actual  analysis  a  precipitate  of  BiOCl  formed  on  diluting  the  solution 
before  passing  in  H2S  in  P.  21  and  that  precipitate  were  filtered  off,  would  a  bis- 
muth precipitate  be  obtained  with  NEUOH  in  P.  35?  if  the  BiOCl  were  not  filtered 
off,  would  any  change  occur  in  it  on  passing  in  H2S? 

14.  Explain  with  the  aid  of  the  mass-action  expressions  involved  why  Cu(OH)2,  a 
substance  very  slightly  soluble  in  water,  is  not  precipitated  by  the  NH4OH.     Show 
that  the  presence  of  the  (NH4)2SC>4  in  the  solution  must  diminish  the  tendency  of  it 
to  precipitate. 

15.  If  the  lead  were  not  removed  by  the  addition  of  £[2804,  would  it  be  precipitated 
as  Pb(OH)2  on  the  addition  of  NEUOH?    What  knowledge  hi  regard  to  lead  com- 
pounds would  enable  one  to  predict  whether  or  not  this  precipitation  would  take 
place? 

16.  Write  the  equations  expressing  the  formation  of  Na2SnO2  from  SnCl2  and 
NaOH;  also  that  expressing  the  spontaneous  decomposition  of  NagSnC^  into  tin 
and  Na2Sn03;  also  that  expressing  its  action  on  BiOgHs. 

17.  Lead  hydroxide,  like  Sn(OH)2,  is  an  amphoteric  substance.    What  is  meant  by 
this  statement?    What  experiments  might  be  made  to  determin  whether  it  is  true? 

18.  If  K4Fe(CN)e  be  added  in  P.  37  to  the  ammoniacal  solution  (without  neutraliz- 
ing it  with  HAc),  no  precipitate  is  produced  unless  a  fairly  large  quantity  of  copper 
is  present.    Explain  this  fact. 

19.  Show  from  the  solubility-product  expressions  that,  if  Cd2Fe(CN)6  is  much  more 
soluble  than  Cu2Fe(CN)6,  the  former  can  not  be  precipitated  till  enough  KiFe(CN)6 
has  been  added  to  precipitate  practically  all  the  copper. 

20.  What  change  is  observed  which  shows  that  the  complex  copper-ammonia  ion 
is  completely  decomposed  when  KCN  is  added  to  the  solution? 

21.  CdS  is  much  more  soluble  in  water  than  Cu2S.    Why  is  CdS  precipitated 
by  H2S  from  the  KCN  solution,  while  Cu2S  is  not? 

22.  In  dissolving  the  sulfides  in  P.  31,  if  a  HNOs  solution  stronger  than  2-normal 
were  used,  or  if  the  solution  were  boiled  too  long,  and  in  consequence  some  mercury 
dissolved,  how  would  it  behave  in  the  subsequent  Procedures? 

23.  If  in  evaporating  the  HNOs  solution  with  H2SC>4  in  P.  33  the  evaporation 
were  stopped  before  all  the  HNOs  was  expelled,  what  effect  would  this  error  have 
on  the  tests  for  lead,  bismuth,  and  cadmium? 

Experiments  8-9. — 1.  Write  chemical  equations  expressing  the  two  stages  of 
the  hydrolysis  of  (NB^S.  Explain  by  the  ionic  theory  and  the  mass-action  law 
why  this  hydrolysis  takes  place,  taking  into  account  the  fact  that  water  is  ionized 
to  a  slight  extent  into  H+  and  OH~. 


QUESTIONS  ON  THE  EXPERIMENTS.  17 

2.  Write  the  equations  expressing  the  action  of  HC1  on  a  solution  of  ammonium 
monosulfide,  and  on  one  of  ammonium  polysulfide  (regarding  the  latter  as  (NH^Sj). 

3.  Write  equations  expressing  the  behavior  of  a  solution  of  AsCla  when  treated 
in  succession  with  H2S  in  P.  21,  with  ammonium  polysulfide  in  P.  22,  and  with  HC1 
in  P.  23. 

4.  Explain  by  the  solubility-product  principle  why  the  addition  of  HC1  to  a  so- 
lution of  (NH^SnSa  causes  the  precipitation  of  81182. 

5.  Make  a  table  showing  the  solubility  of  each  of  the  sulfides  of  the  copper  and 
tin  groups  in  both  ammonium  monosulfide  and  ammonium  polysulfide.     Indicate 
in  each  case  whether  the  sulfide  is  readily  soluble,  slightly  soluble,  or  practically 
insoluble. 

6.  Why  does  not  Bi2S3  dissolve  in  (NH^S,  just  as  Sb2S3  does?    Why  does  not 
SnS  dissolve  in  ammonium  monosulfide  as  well  as  in  ammonium  polysulfide? 

7.  Why  not  use  in  all  cases  ammonium  polysulfide,  since  this  readily  dissolves  all 
the  tin-group  sulfides?     Why  is  the  precipitate  first  treated  with  a  small  quantity 
of  it,  even  tho  this  makes  necessary  a  second  treatment? 

8.  Why  does  the  tendency  of  copper  to  form  the  complex  copper-ammonia  ion 
not  cause  CuS  to  dissolve  in  NEUOH?     What  must  be  the  explanation  of  the  fact 
that  some  CuS  dissolves  in  ammonium  polysulfide? 

Experiment  10. — 1.  Describe  specifically  the  differences  in  the  behavior  of  the 
sulfides  of  arsenic,  antimony,  and  tin  on  which  the  separation  of  these  three  ele- 
ments is  based. 

2.  In  treating  the  sulfides  with  12-normal  HC1  why  does  much  more  As2S5  dis- 
solve if  the  solution  be  allowed  to  boil?    Why  does  it  boil  at  so  low  a  temperature 
as  50-60°? 

3.  Write  by  the  method  described  in  Note  7,  P.  21,  the  equations  expressing  the 
action  of  HC1  on  KClOs  by  which  C\z  is  produced  and  that  by  which  C1C>2  is  pro- 
duced. 

4.  Explain  why  the  Ck  set  free  by  the  addition  of  the  KClOs  causes  the  As2Ss 
to  dissolve  even  in  the  dilute  HC1. 

5.  What  is  the  expression  for  the  solubility-product  in  the  case  of  MgNH4As04? 
Why  does  it  dissolve  readily  in  HC1? 

6.  Why  does  the  hydrolysis  of  this  salt  increase  its  solubility?    Why  is  that 
hydrolysis  decreased  by  an  excess  of  NH^OH?     How  is  the  hydrolysis  affected  by 
the  presence  of  NH4C1?    Would  NH4C1  affect  the  solubility  in  any  other  way? 

7.  What  is  a  saturated  solution?  a  supersaturated  one?     By  what  treatments  can 
a  precipitate  be  made  to  separate  from  a  supersaturated  solution? 

8.  What  difference  in  the  ionization  causes  the  behavior  of  arsenic  acid  towards 
E^S  to  be  different  from  that  of  other  elements  of  the  copper  and  tin  groups?    What 
causes  its  own  behavior  to  be  different  in  dilute  and  concentrated  HC1  solutions? 

9.  Write  the  series  of  reactions  which  take  place  when  H^S  is  passed  into  a  dilute 
HC1  solution  of  HgAsOi.      State  what  is  known  about  the  rate  or  equilibrium  of 
each  of  these  reactions. 

10.  Explain  why  antimony  precipitates  on  the  platinum  rather  than  on  the  tin, 
and  state  how  would  the  result  be  different  if  the  tin  did  not  touch  the  platinum? 

11.  Why  does  the  tin  dissolve  in  HC1  more  readily  when  it  is  in  contact  with  the 
platinum? 


18  QUESTIONS  ON  THE  EXPERIMENTS. 

12.  In  the  confirmatory  test  for  tin,  why  is  lead  used  instead  of  zinc? 

13.  If  zinc  were  used,  how  would  the  procedure  have  to  be  modified? 

14.  In  the  confirmatory  test  for  tin  how  does  the  precipitation  of  a  mercury  com- 
pound show  the  presence  of  tin? 

15.  If  the  antimony  were  not  all  precipitated  in  P.  44  on  passing  H2S  into  the  hot 
solution,  how  would  it  behave  in  the  subsequent  procedure  (P.  46)  in  which  the 
tin  is  precipitated  and  its  presence  confirmed? 

Experiment  12. — 1.  In  an  actual  analysis  how  many  cubic  centimeters  of  NH4OH 
would  be  required  to  neutralize  the  5  cc.  HNOs  that  are  added  before  precipitating 
with  H2S? 

2.  How  much  more  NH4OH  would  be  needed  to  neutralize  the  solution  if  500  mg. 
Cu  had  been  present  in  the  form  of  Cu(NOs)2  in  the  solution  precipitated  by  H2S? 
(In  all  such  calculations  of  the  volume  of  the  reagent  needed,  first  reduce  the 
weight  of  the  constituent  from  grams  to  equivalents.) 

3.  How  does  testing  the  vapors  above  the  solution  with  PbAc2-paper  show  that 
an  excess  of  (NH^S,  a  non-volatil  salt,  has  been  added? 

4.  If  in  an  actual  analysis  the  mixture  containing  NH4OH  and  (NH4)2S  were 
allowed  to  absorb  CC>2  from  the  air  before  filtering,  what  difference  would  it  make? 

5.  Why  is  the  (NHi)2S  precipitate  treated  first  with  cold  HC1?    Why  is  HNO3 
subsequently  added? 

Experiments  13  and  14. — 1.  Which  elements  are  soluble,  a,  in  excess  of  NH4OH 
(in  the  presence  of  NH4C1),  but  not  in  excess  of  NaOH;  6,  in  excess  of  NaOH,  but 
not  of  NH4OH  (in  presence  of  NH4C1);  c,  in  excess  both  of  NBUOH  and  of  NaOH; 
d,  neither  in  excess  of  NaOH  nor  of  NH4OH  (hi  presence  of  NHjCl)? 

2.  What  are  the  explanations  of  the  four  typical  cases  a,  6,  c,  d,  referred  to  in  the 
preceding  question? 

3.  Could  the  hydroxide  of  an  element  which  does  not  form  a  complex  ammonia 
cation  be  soluble  in  NHiOH  and  not  in  NaOH?    Could  an  amphoteric  hydroxide 
be  readily  soluble  in  NaOH  and  entirely  insoluble  in  NH4OH? 

4.  Show  by  formulating  and  combining  the  two  mass-action  equations  involved 
that  the  quantity  of  aluminum  dissolved  (as  A1O2~)  in  the  presence  of  a  base  is  pro- 
portional to  the  OH~  concentration  in  the  solution. 

5.  Name  all  the  elements  that  form  ammonia  complexes  in  all  the  groups  thus 
far  considered.    What  can  be  said  as  to  the  position  of  these  elements  in  the  periodic 
system?     (Refer  to  a  text-book  of  Inorganic  Chemistry.) 

6.  If  in  an  actual  analysis  no  precipitate  is  obtained  on  the  addition  of  NEUOH, 
what  conclusion  may  be  drawn? 

7.  Which  of  the  hydroxides  precipitated  by  NaOH  undergo  change  on  the  addi- 
tion of  NaaO2,  and  into  what  compound  is  each  of  these  hydroxides  converted? 

8.  What  substances  are  produced  by  the  action  of  Na^Oa  on  water? 

Experiment  15. — 1.  What  does  Expt.  15  show  in  regard  to  the  behavior  of  the 
alkaline-earth  elements  in  an  actual  analysis? 

2.  What  must  be  the  explanation  of  the  fact  that  the  phosphate  combines  with 
the  iron  rather  than  the  calcium  when  both  these  elements  are  present? 

3.  If  phosphate  is  known  to  be  present,  is  it  necessary  to  test  for  alkaline-earth 
elements  in  the  filtrate  from  the  (NH4)2S  precipitate? 

4.  If  the  original  substance  were  soluble  in  water,  would  it  be  necessary  to  test 
for  these  elements  in  the  (NH4)2S  precipitate? 


QUESTIONS  ON  THE  EXPERIMENTS.  19 

5.  If  CaCO3  were  substituted  for  Ca^(POi)2  in  the  first  part  of  Expt.  15,  what 
would  the  result  have  been? 

6.  If  a  solution  of  CaaCPO^  in  HC1  were  treated  by  the  second  paragraph  of 
P.  52,  what  would  happen  on  the  addition  of  each  reagent?    What  would  happen, 
if  a  large  proportion  of  Feds  were  also  present  in  the  phosphate  solution? 

Experiment  16. — 1.  Make  a  table  (like  that  described  in  Question  1  on  Expt.  6) 
showing  the  operations  involved  and  the  behavior  of  the  chromium  and  zinc  in 
analyzing  a  dilute  HNOs  solution  of  ZnCrO^  beginning  with  the  K^S  precipitation 
(P.  21)  and  continuing  through  the  analysis  of  the  aluminum-group  (P.  51-57). 
At  the  foot  of  the  table  write  all  the  chemical  equations  involved. 

2.  In  separating  the  aluminum  from  the  chromium  and  zinc  with  NRiOH  in  P. 
53,  what  would  be  the  harm  of  adding  too  small  an  excess?    what  of  adding  too 
large  an  excess? 

3.  By  what  reagent  other  than  BaC^  could  the  chromate  be  precipitated?    What 
disadvantage  would  there  be  in  the  use  of  this  reagent? 

4.  How  can  sulfate  be  present  in  the  solution  to  which  BaCl2  is  added? 

5.  How  could  a  precipitate  of  ZnS  be  distinguished  from  one  of  sulfur  by  adding 
a  suitable  reagent  to  the  liquid  containing  it? 

6.  What  happens  to  nitrates,  such  as  Zn(NC>3)2  or  Co(NOs)2,  when  they  are  ignited, 
as  in  the  confirmatory  test  for  zinc? 

Experiment  17. — 1.  What  are  the  oxides  of  manganese  corresponding  to  its  three 
stages  of  oxidation  occurring  in  P.  61  and  62?  What  is  the  valence  of  manganese 
in  each  of  these  oxides?  How  do  they  differ  with  respect  to  the  formation  of  salts 
with  acids  and  with  bases? 

2.  Make  a  table  (like  that  described  in  Question  1  on  Expt.  6)  showing  the 
operations  involved  and  the  behavior  of  manganese  in  analyzing  a  dilute  HNOs 
solution  of  CaMnO*  beginning  with  the  H2S  precipitation  and  continuing  through 
the  final  test  for  manganese  (thus  involving  P.  21,  51,  52,  61,  and  62).     Write  also 
all  the  chemical  equations  involved. 

3.  Why  is  a  considerable  excess  of  NH4OH  added  in  precipitating  the  iron  in 
P.  64? 

4.  If  FeS  were  treated  by  P.  31-38,  how  would  it  behave  with  each  of  the  reagents? 

5.  Why  is  it  necessary  to  test  for  zinc  in  the  analysis  of  the  iron-group? 

6.  Why  may  zinc  be  precipitated  by  NaOH  and  Na2O2  in  the  first  treatment 
(in  P.  52),  and  yet  not  be  precipitated  by  them  in  the  second  treatment  (in  P.  67)? 

7.  When  the  original  Na2O2  precipitate  is  so  small  that  it  need  not  be  tested  for 
zinc,  how  may  P.  67-68  be  simplified? 

8.  When  the  H2§  precipitate  obtained  in  P.  66  is  small,  how  may  P.  67-68  be 
simplified? 

Experiment  18. — 1.  If  the  H2S  and  NaA  precipitates  were  not  washed  free  from 
ammonium  and  sodium  salts  before  dissolving  them  in  P.  68,  what  might  happen 
upon  the  addition  of  the  HCl-ether  reagent? 

2.  Suggest  an  explanation  of  the  fact  that,  tho  NiCl2  is  precipitated  as  a  yellow 
compound  by  the  HCl-ether  reagent,  it  has  a  green  color  even  in  very  concentrated 
aqueous  solution. 

3.  Suggest  any  possible  differences  in  the  molecular  state  of  CoCk,  which  might 
account  for  the  facts  that  it  has  a  pink  color,  like  that  of  Co(NO3)2,  in  a  dilute 
aqueous  solution,  and  a  blue  color  in  an  ether  solution  or  a  concentrated  aqueous 
solution  saturated  with  HC1. 


20  QUESTIONS  ON  THE  EXPERIMENTS. 

4.  Name  all  the  elements  thus  far  considered  which  in  any  state  of  oxidation  form 
colored  compounds  in  solution.     What  can  be  said  as  to  the  position  of  these 
elements  in  the  periodic  system? 

5.  Write  equations  expressing  the  steps  in  the  chemical  process  by  which  the  pre- 
cipitate of  potassium  cobaltinitrite  may  be  considered  to  be  formed. 

6.  Does  the  N(V  coming  from  the  excess  of  KNO2  diminish  the  solubility  of  the 
precipitate  in  the  same  way  as  the  K+  does?    What  else,  from  a  mass-action  stand- 
point, might  it  be  expected  to  do? 

Experiment  19. — 1.  Why  must  the  HNOs  be  removed  before  testing  for  iron  with 
KSCN? 

2.  Make  a  table  showing  the  operations  involved  and  the  behavior  of  Ca3(PO4)2 
in  analyzing  a  HNO3  solution  of  it  by  P.  51-52,  P.  61,  and  P.  65. 

3.  Why  does  the  NH4Ac  solution  become  red  only  when  the  quantity  of  iron  is 
more  than  equivalent  to  the  quantity  of  phosphate  present? 

4.  What  is  meant  by  a  basic  salt?    What  are  two  possible  simple  formulas  for 
basic  ferric  acetate? 

Experiment  21. — 1.  What  does  this  experiment  show  as  to  the  precipitation  by 
(NH^COs  of  magnesium  and  of  the  other  alkaline-earth  elements  (which  all  behave 
nearly  as  calcium  does)? 

2.  Why  would  a  reagent  which  consisted  of  NH^COs  not  be  suitable  for  the 
precipitation? 

3.  Why  is  there  any  advantage  in  adding  more  NHs  than  corresponds  to  the  neu- 
tral salt  (NH4)2CO3? 

4.  If  it  were  desired  to  work  out  a  procedure  for  separating  calcium,  barium,  and 
strontium  from  magnesium  by  means  of  (NH  ^COs,  what  experiments  would  one 
naturally  make? 

Experiment  22. — 1.  In  order  to  make  a  separation  of  1  mg.  barium  from  500  mg. 
strontium,  what  must  be  the  concentration  of  CrO4~,  stated  with  reference  to  the 
saturation-values  of  the  ion-concentration  products  of  BaCrO4  and  SrCrO4? 

2.  What  must  be  true  of  the  relativ  values  of  these  two  saturation-values  in  order 
that  this  separation  may  be  possible?     What  is  the  actual  ratio  of  these  two  values? 
(See  the  Table  of  Solubilities  on  page  124.) 

3.  Write  the  chemical  equations  for  the  conversion  of  chromate-iron  into  hydro- 
chromate-ion,  and  for  the  conversion  of  the  latter  into  bichromate-ion.    Write  also 
the  mass-action  expressions  for  the  equilibrium  of  these  reactions.    Show  by  them 
what  determins  the  proportion  of  CrC>4=  and  of  HCrCU"  in  any  solution,  and  what 
determins  the  concentration  of  C^O?"  in  any  solution. 

4.  In  practis  what  three  substances  must  be  added  in  proper  proportions  in  order 
to  secure  the  right  CrC>4~  concentration  in  the  solution? 

5.  Why  is  the  second  K2CrO4  precipitate  obtained  in  the  confirmatory  test  for 
barium  more  conclusiv  evidence  of  its  presence  than  the  first  K2CrO4  precipitate? 

6.  On  addition  of  NEUOH  in  P.  84  what  chemical  change  causes  the  change  in 
color  from  orange  to  yellow?    Why  from  a  mass-action  standpoint  does  the  addition 
of  NEUOH  cause  this  change  to  take  place?    Why  does  this  change  cause  strontium 
to  precipitate? 

7.  Explain  fully  with  reference  to  the  saturation-values  of  the  ion-concentration 
products  why  the  carbonate-chromate  mixture  used  in  P.  85  does  not  affect  BaCrC>4, 
and  why  it  converts  SrCrO4  into  SrCO3. 


QUESTIONS  ON  THE  EXPERIMENTS.  21 

8.  Why  does  CaQjCU  dissolve  in  dilute  H2SO4,  but  not  in  HAc? 

9.  How  does  the  confirmatory  test  for  calcium  distinguish  it  from  barium  and 
strontium,  which  form  much  less  soluble  sulfates?    How  does  it  distinguish  calcium 
from  magnesium? 

10.  Could  magnesium  be  precipitated  by  any  other  reagent  in  the  form  of  a  com- 
pound closely  analogous  to  magnesium  ammonium  phosphate? 

11.  Why  is  the  production  of  a  precipitate  with  NaaHPOi  in  the  confirmatory 
test  for  magnesium  (in  P.  89)  more  conclusiv  evidence  of  its  presence  than  the 
production  of  the  first  Na^HPC^  precipitate  (in  P.  88)? 

Experiment  23. — 1.  If  the  ammonium  salt  were  not  completely  removed  by  the 
ignition,  how  would  it  behave  in  the  subsequent  tests  for  potassium? 

2.  Why  is  the  separation  of  potassium  and  sodium  by  the  HC1O4  method  satis- 
factory when  these  elements  are  present  as  chlorides  or  nitrates,  but  not  when  they 
are  present  as  sulfates?    Why  is  it  satisfactory  when  they  are  present  as  phos- 
phates? 

3.  What  is  implied  by  the  statement  that  Na3Co(NO2)e  is  a  complex  salt  in 
solution? 

4.  How  might  a  solution  of  NaaCotNQOe  be  prepared,  judging  from  previous  expe- 
rience with  an  analogous  compound? 

Experiments  26  and  27. — 1.  Describe  in  detail  the  method  by  which  it  can  be 
shown  that  the  magnetic  oxide  of  iron  (FesOJ  contains  both  ferrous  and  ferric  iron. 

2.  Describe  a  method  by  which  it  can  be  shown  that  a  solution  contains  both 
mercurous  and  mercuric  salts. 

3.  What  indications  of  the  state  of  oxidation  of  tin  would  be  afforded  by  the 
treatment  with  H2S  in  P.  21,  and  with  (NH4)2S  in  P.  22,  of  a  solution  of  SnCl2?  of 
solution  of  SnCl4? 

4.  What  other  constituents  besides  iron,  mercury,  and  tin  exist  in  two  stages  of 
oxidation  in  the  form  of  salts?    Giv  the  symbols  of  compounds  illustrating  these 
two  states  of  oxidation. 

5.  Mention  any  phenomena  that  might  be  observed  during  the  course  of  analysis 
which  would  distinguish  from  each  other  the  two  stages  of  oxidation  of  each  of  the 
constituents  named  in  the  answer  to  Question  4. 

Experiments  28  and  29. — 1.  What  indication  of  the  state  of  oxidation  of  arsenic 
may  be  obtained  in  the  course  of  the  analysis  for  the  basic  constituents? 

2.  What  do  the  facts  that  concentrated  HC1  converts  much  H3AsO3,  but  very 
little  HsAsO4,  into  the  corresponding  chloride  show  as  to  the  ionization  of  these  two 
substances? 

3.  What  reaction  occurs  when  the  CuSO4  solution  is  poured  into  the  hydrogen 
generator  in  P.  99?    How  does  the  addition  of  the  CuSO4  accelerate  the  evolution 
of  hydrogen? 


22  QUESTIONS   ON    THE   EXPERIMENTS. 

DETECTION    OF   THE   ACIDIC    CONSTITUENTS. 

Experiment  31. — 1.  With  the  aid  of  the  laboratory  experience  and  the  statement 
as  to  solubilities  on  page  31  arrange  the  barium  salts  of  the  acidic  constituents  listed 
on  page  93  in  four  groups  comprising  respectivly:  those  readily  soluble  in  water; 
those  only  slightly  soluble  in  water,  but  readily  soluble  in  HAc;  those  only  slightly 
soluble  in  HAc,  but  readily  soluble  in  dilute  HC1;  and  those  only  slightly  soluble  in 
dilute  HC1. 

2.  Explain  the  following  facts  with  reference  to  the  ionization-values  given  in  the 
Table  on  page  123:  a,  a  precipitate  of  BasfPO^  is  converted  into  one  of  BaHPO4 
by  a  nearly  equivalent  quantity  of  HAc;  6,  BaHPO4  is  readily  soluble  in  an  excess  of 
HAc  with  formation  of  Ba(H2PC>4)2  and  BaAc2;  c,  BaHPCU  is  readily  soluble  in  dilute 
HC1  with  formation  of  H3PC>4  and  BaCl2,'  d,  BaSO3  is  not  much  more  soluble  in 
HAc  than  in  water,  but  is  readily  soluble  in  dilute  HC1;  e,  BaCrO4  is  not  readily 
soluble  in  HAc,  even  tho  HCrO4~  is  very  slightly  ionized. 

3.  State  how  fluoride  when  present  in  large  quantity  behaves  in  each  step  of  P. 
102;  and  explain  this  behavior  with  reference  to  the  ionization  of  HF  and  the  solu- 
bility-values given  in  the  table  on  page  124. 

4.  What  other  acidic  constituent  besides  fluoride  is  shown  by  the  solubility-table 
on  page  124  to  form  a  calcium  salt  less  soluble  than  the  barium  salt?     Show  by 
reference  to  the  solubility  and    ioniaation  values  involved  how  this   constituent 
(in  comparison  with  fluoride)  might  be  expected  to  behave  in  each  step  of  P.  102. 

5.  Suggest  a  reason  why  chromate,  tho  included  in  P.  102,  is  not  added  to  the  mix- 
ture (of  Na2S04,  Na2SO3,  and  KF)  used  in  Expt.  31  in  illustrating  that  Procedure. 

Experiment  32. — 1.  What  disadvantage  would  there  be  in  substituting  Pb(NO3)2 
for  Cd(NO3)2  in  the  test  for  sulfide  in  P.  103? 

2.  Arrange  the  silver  salts  of  the  acidic  constituents  listed  on  page  93  in  three 
groups  comprising  respectivly:  those  readily  soluble  in  water;  those  only  slightly 
soluble  in  water,  but  readily  soluble  in  dilute  HNO3;  and  those  only  slightly  soluble 
in  dilute  HNO3. 

3.  Explain  by  reference  to  the  solubilities  of  the  salts  (as  given  in  the  Table  on 
page  124)  and  the  ionizations  of  the  corresponding  acids  why  each  of  the  salts  in  the 
second  group  referred  to  in  the  previous  question  dissolves  in  dilute  HNO3. 

4.  Explain  why  the  silver  halides  do  not  dissolve  in  dilute  HNO3. 

5.  Explain  why  Ag2S  is  only  slightly  soluble  in  dilute  HNO3,  even  tho  the  ionization 
of  HS~  is  extremely  small. 

6.  Explain  why  Ag2(CN)2  is  only  slightly  soluble  in  dilute  HNO3,  even  tho  HCN 
is  only  very  slightly  ionized. 

7.  Why  would  it  not  be  satisfactory  to  precipitate  (and  thus  partially  identify  by 
their  color)  the  silver  salts  which  are  soluble  in  dilute  HNO3,  but  not  in  water,  by 
adding  to  the  filtrate  from  the  silver  halides  an  excess  of  NaOH?     Why  not  by  adding 
an  excess  of  NH4OH? 

Experiment  35. — 1.  State  the  principles  involved  in  the  process  used  for  the  detec- 
tion of  the  three  halides  in  the  presence  of  each  other. 

2.  Does  the  fact  that  only  a  very  small  quantity  of  bromin  is  liberated  in  the  first 
step  of  the  process  mean  that  the  reaction  between  the  bromide  and  KMnC>4  is  in 
equilibrium  under  the  conditions  prevailing  in  the  solution? 

3.  What  is  meant  by  the  statement  that  a  reaction  is  in  equilibrium?     How  would 
one  proceed  to  determin  whether  a  given  reaction-mixture  is  in  equilibrium? 


QUESTIONS  ON  THE  EXPERIMENTS.  23 

4.  If  the  mixture  of  KBr,  NaAc,  HAc,  and  KMn04  were  allowed  to  stand  a  day 
or  a  week,  what  would  happen? 

5.  Why  must  the  iodin  set  free  in  the  first  part  of  the  process  be  completely 
extracted  by  the  chloroform? 

6.  Why  is  H^SOs  added  in  the  last  part  of  the  process?    Write  the  equation  for 
the  reaction  which  it  causes. 

7.  Why  is  HNO3  added  with  the  AgNO3  at  the  end  of  the  process? 

Experiment  37. — 1.  What  two  things  detennin  whether  or  not  an  acid  passes  over 
into  the  first  distillate? 

2.  Show  by  reference  to  the  mass-action  expressions  for  the  ionization  of  the  two 
acids  what  deterrnins  the  extent  to  which  an  acid  is  displaced  from  its  salt  by  another 
acid,  taking  K+CN~  and  H+H2PO4~  as  an  example. 

3.  Explain  by  reference  to  the  ionization- values  involved  why  BaCOs  dissolves 
on  adding  HAc,  and  why  BaSOs  does  not. 

4.  Show  how  phosphoric,  pyrophosphoric,  and  metaphosphoric  acids  are  related 
to  one  another  in  composition. 

5.  If  a  compound  of  an  element  forming  an  insoluble  phosphate  (for  example, 
CaCOs)  were  distilled,  with  HsPO^  would  the  insoluble  phosphate  separate  in  the 
distilling  flask?     Explain  why  or  why  not. 

6.  If  E^SOs  is  found  in  the  first  distillate,  which  of  the  other  substances  that  may 
be  in  that  distillate  is  it  unnecessary  to  test  for?    Write  the  equation  for  the  reac- 
tion which  would  take  place  between  H^SOs  and  each  of  these  substances. 

7.  If  any  of  the  HsPC^  were  thrown  over  mechanically  into  the  first  distillate, 
what  erroneous  conclusion  might  be  drawn  in  connection  with  P.  111? 

Experiment  38. — 1.  What  different  constituents  of  the  original  substance  may 
giv  rise  to  sulfur  in  the  distillate?  what  ones  to  sulfurous  acid? 

2.  Show  that  HsPC^  which  had  been  thrown  over  mechanically  into  the  first  dis- 
tillate would  not  interfere  with  the  tests  for  sulfite  and  carbonate  in  P.  112. 

Experiment  39. — 1.  What  different  constituents  in  the  original  substance  may 
giv  rise  to  chlorin  in  the  first  distillate?  Write  chemical  equations  illustrating  its 
production  from  each  of  these  constituents. 

2.  What  different  substances  that  might  be  present  in  the  distillate  would  cause 
iodin  to  be  set  free  on  the  addition  of  KI? 

3.  What  conclusions  could  be  drawn  from  the  tests  of  the  first  two  paragraphs 
of  P.  113  as  to  the  presence  or  absence  of  each  of  the  three  halogens  in  a  distillate 
containing  the  following  halogens:  a,  12,  Br2,  and  Cfe;  b,  Br2j  c,  Ck;  d,  no  halogen? 

4.  Explain  the  fact  that  iodin  is  extracted  by  chloroform  more  slowly  from  an 
aqueous  solution  when  it  contains  iodide.     State  the  law  that  determins  the  quan- 
tity extracted. 

5.  Write  chemical  equations  illustrating  the  process  by  which  a  small  quantity  of 
HNO2  liberates  a  large  quantity  of  iodin  from  KI. 

Experiment  40. — 1.  Write  the  chemical  equations  involved  in  the  test  for  cyanide. 

2.  Into  what  compounds  is  K4Fe(CN)e  decomposed  when  it  is  distilled  with  HaPC^? 

3.  Referring  to  the  results  of  Expt.  26  with  ferrous  and  ferric  salts,  suggest  how 
a  ferrocyanide  could  be  distinguished  from  a  ferricyanide  by  tests  applied  to  a  solu- 
tion of  the  original  substance. 

Experiment  41. — 1.  To  what  extent  is  the  analysis  of  the  second  distillate  sim- 
plified when  AgNOa  givs  no  precipitate? 


24  QUESTIONS  ON  THE  EXPERIMENTS. 

2.  Some  H3PC>4  may  pass  over  into  the  second  distillate.    Would  Ag3PO4,  which 
is  only  slightly  soluble  in  water,  precipitate  in  the  AgNO3  test,  and  thus  obscure  the 
test  for  the  other  constituents? 

3.  When  is  it  necessary  to  test  the  second  distillate  for  sulfide? 

4.  Explain  by  the  mass-action  principles  why  a  sulfide  may  giv  off  an  appreciable 
quantity  of  H2S  only  in  the  second  part  of  the  distillation. 

5.  What  constituents  in  the  substance  are  likely  to  giv  rise  to  chlorin  in  the  second 
distillate?  to  bromin?  to  iodin? 

Experiment  42. — 1.  Why  not  test  for  boric  acid  in  the  distillates  obtained  in  the 
H3PO4  distillation? 

2.  What  is  the  advantage  in  P.  121  of  distilling  the  borate  with  H2SC>4  and  CHaOH, 
rather  than  with  H2S04  with  water? 

Experiment  43. — 1.  Why  must  all  the  substances  used  in  the  test  for  fluoride  be 
thoroly  dry? 

2.  What  happens  to  KHS04  when  it  is  heated? 

3.  What  are  the  main  constituents  of  glass?    What  is  the  action  of  HF  on  it? 

4.  Since  H2Si03  is  non-volatil,  how  can  the  deposit  be  driven  up  the  tube  by 
heating? 

5.  Of  what  does  the  white  deposit  left  after  washing  the  tube  with  water  consist? 

6.  Show  why  it  is  appropriate  to  call  the  compound  H2SiFe  "fluo-silicic  "  acid. 

7.  Why  is  it  not  satisfactory  to  test  for  fluoride  in  the  distillates  obtained  in  the 
HsP04  distillation? 

Experiments  44  and  45. — 1.  If  chlorate  were  present  in  the  substance,  what  would 
happen  to  it  when  treated  by  P.  124? 

2.  Iodin  is  not  completely  reduced  to  HI  by  FeS04,  the  four  substances  I2,  HI, 
ferrous  salt,  and  ferric  salt  all  being  present  in  considerable  quantity  at  equilibrium. 
If  a  solution  containing  ferric  sulfate  and  KI  were  submitted  to  P.  124,  what  would 
be  the  result? 

3.  How  might  the  distillation-process  of  P.  124  be  modified  so  as  to  detect  in 
two  steps,  first,  the  presence  of  nitrite,  and  second,  when  it  is  absent,  the  presence 
of  nitrate? 

Experiment  46. — 1.  What  different  substances  are  present  in  the  solution  pro- 
duced by  acidifying  the  NaOCl  solution  with  HAc? 

2.  What  advantage  would  there  be  in  making  the  hypochlorite  test  in  alkaline 
solution,  rather  than  in  HAc  solution?    What  disadvantage? 

3.  Assuming  that  the  oxidation  of  the  PbAc2  to  PbO2  is  caused  by  HOC1,  but  not 
by  Ck,  explain  why  the  oxidation  does  not  take  place  in  a  solution  strongly  acidified 
with  HNO3. 

Experiment  47. — 1.  How  may  a  chlorate  giv  off  chlorin  in  the  first  part  of  the 
HsPO4  distillation?  How  in  the  second  part? 

2.  Write  the  equation  expressing  the  reaction  that  takes  place  between  NaOCl  and 
NaAsO2. 

3.  If  the  hypochlorite  were  not  reduced,  AgCl  and  AgClO3  would  be  formed  in 
the  HNO3  solution  by  the  action  of  the  C12  on  the  AgN03.    Write  the  equation  ex- 
pressing this  reaction. 


QUESTIONS  ON  THE  EXPERIMENTS.  25 


PREPARATION    OF   THE    SOLUTION. 

Experiment  48. — 1.  Of  what  elements  is  organic  matter  mainly  composed,  and 
what  causes  it  to  blacken  on  heating? 

2.  Describe  a  method  by  which  a  substance  might  be  tested  for  organic  matter 
by  converting  the  carbon  into  carbon  dioxide. 

3.  Could  a  sirup  be  tested  for  aluminum  by  diluting  it  with  water  and  analyzing 
the  solution  in  the  usual  way?    Could  it  be  so  tested  for  iron? 

4.  Name  the  different  forms  in  which  water  may  be  present  in  a  substance. 

5.  If  a  substance  were  ignited  at  a  red  heat  (for  example,  to  destroy  organic 
matter)  before  submitting  it  to  analysis,  what  basic  constituents  would  be  lost? 

6.  State  how  each  of  the  following  substances  would  behave  in  the  closed-tube 
test:  a,  Na3AsO4;  b,  MgNH4PO4;  c,  FeS2;  d,  Pb(NO3)2;  e,  NaC2H3O2. 

7.  What  acidic  constituents  will  be  detected,  when  present  in  considerable  quan- 
tity, upon  treating  the  substance  with  dilute  HNO3? 

8.  What  is  the  behavior  towards  litmus  of  an  aqueous  solution  of  each  of  the 
following  substances:  Na2SO4;  Na^COs;  NaHCO3;  NaHSO4;  Ca(NO3)2;  Fe(NO3)3? 

9.  Name  the  basic  constituents  whose  salts  are  nearly  all  readily  soluble  in  water. 

10.  Name  the  basic  constituents  whose  salts  are  decomposed  by  water,  but  dis- 
solved by  dilute  acids. 

11.  Name  the  acidic  constituents  whose  salts  are  all  or  nearly  all  readily  soluble 
in  water. 

12.  Name  the  acidic  constituents  whose  salts  (except  those  of  the  alkali  elements) 
are  nearly  all  very  slightly  soluble  in  water. 

13.  Of  the  groups  of  salts  named  in  the  answer  to  Question  12,  which  are  readily 
soluble  in  cold  dilute  HNO3  or  HC1?    Why? 

14.  In  a  substance  soluble  in  water  which  has  been  found  to  contain  barium,  which 
of  the  following  constituents  would  it  be  unnecessary  to  test  for:  nitrate,  phosphate, 
sulfide,  sulfite,  chloride,  sulfate,  carbonate? 

Experiments  49  and  50. — 1.  What  group  of  salts  not  dissolved  by  cold  dilute 
HNO3  are  decomposed  by  hot  concentrated  HNO3,  because  of  its  oxidizing  action? 
what  groups  because  of  its  action  as  a  strong  acid? 

2.  State  what  happens  (that  is,  what  chemical  changes  occur  and  what  phenomena 
are  observed)  at  each  step  of  the  process  when  PbSiO3  (which  is  decomposed  by 
strong  acids)  is  treated  by  P.  3. 

3.  State  what  happens  at  each  step,  namely,  on  heating  with  strong  HNO3,  on 
evaporating  to  dryness  and  adding  cold  dilute  HNO3,  and  on  heating  the  residue 
with  strong  HC1,  when  each  of  the  following  substances  is  treated  by  P.  3  and  4: 
a,  Sb2S3;  6,  MnO2;  c,  PbSO4. 

4.  What  metallic  elements  are  scarcely  ever  found  in  alloys?     How  may  the 
process  of  analysis  therefore  be  shortened? 

5.  If  there  is  no  residue  after  treating  an  alloy  by  P.  3,  what  does  it  show  as  to 
the  absence  of  certain  elements? 

Experiment  51. — 1.  State  what  happens  to  each  of  the  substances  whose  sym- 
bols are  given  under  b  in  Table  I  on  fusing  it  with  Na2CO3,  on  treating  the  fused 
mass  with  water,  and  on  treating  the  residue  with  dilute  HNO3,  as  described 
in  P.  7. 

2.  What  are  the  only  acidic  constituents  that  it  is  ordinarily  necessary  to  test  for 
in  minerals? 


26  QUESTIONS  ON  THE  EXPERIMENTS. 

3.  State  what  happens  to  each  of  the  substances  whose  symbols  are  given  under 
6  in  Table  I  when  it  is  treated,  as  in  P.  5:  (1)  with  concentrated  H^SO^*  (2)  then 
with  HF;  (3)  then,  after  evaporating  to  fuming,  with  water. 

4.  State  what  happens  to  feldspar  (potassium  aluminum  silicate,  an  example  of 
a  silicate  not  much  decomposed  by  acids  other  than  HF)  when  it  is  treated  by  P.  5. 

5.  Explain  with  reference  to  the  solubilities  of  the  substances  involved  why 
PbS(>4  is  completely  converted  into  PbCOs  by  boiling  with  Na^COa  solution  in 
P.  6;  and  why  BaSC>4,  when  much  of  it  is  present,  is  only  partly  converted  into 
BaCOs  by  the  same  treatment. 


PART  II. 
THE  SYSTEM  OF  ANALYSIS. 


PREPARATION  OF  THE  SOLUTION. 


PRELIMINARY    EXAMINATION. 

Procedure  i. — Preliminary  Examination. — If  the  substance  is  a 
non-metallic  solid,  note  its  color,  odor,  and  texture;  examin  it  with  a 
lens  to  determin  whether  it  is  heterogeneous,  and,  if  so,  note  the  appear- 
ance of  its  constituents.  To  determin  whether  organic  matter  or 
water  is  present  and  to  get  other  indications,  heat  gently  at  first, 
then  strongly,  about  0.1  g.  (0.1  gram)  of  the  finely  powdered  substance 
in  a  hard  glass  tube  (of  about  0.6  cm.  bore  and  10  cm.  length)  closed 
at  one  end.  Note  whether  the  substance  blackens,  whether  a  tarry, 
aqueous,  or  other  deposit  forms  on  the  cold  part  of  the  tube,  and 
whether  any  odor  is  emitted.  If  organic  matter  is  thus  proved  to  be 
absent,  pass  to  P.  2  (Procedure  2) ;  if  proved  to  be  present,  to  P.  8. 

If  the  substance  is  an  alloy,  treat  it  by  P.  3. 

If  the  substance  is  a  solution,  evaporate  a  measured  volume  of  it 
to  dryness  in  a  small  weighed  dish,  dry  the  residue  thoroly  at  120- 
130°  in  a  hot  closet  or  by  keeping  the  dish  in  motion  over  a  small 
flame,  and  weigh  the  dish  again.  Heat  a  portion  of  this  residue  in  a 
closed  tube  as  described  above.  Treat  another  portion  by  P.  2  if 
organic  matter  is  absent,  or  by  P.  8  if  organic  matter  is  present. 

Notes. — 1.  When  a  complete  analysis  in  the  wet  way  is  to  be  made,  it  is 
usually  not  worth  while  to  make  a  more  extended  preliminary  examination  in 
the  dry  way.  The  closed-tube  test  is,  however,  essential,  in  order  to  show 
whether  organic  matter  is  present;  for  certain  kinds  of  organic  matter, 
especially  sugars  and  hydroxy-acids,  such  as  tartaric,  citric,  and  lactic  acids, 
prevent  the  precipitation  of  the  hydroxides  of  most  of  the  elements  by  alkalies. 
Such  organic  matter  must  therefore  be  detected  and  removed  in  order  to  ensure 
the  precipitation  of  aluminum  and  chromium  by  NEUOH.  Moreover,  a  large 
quantity  of  organic  matter  of  any  kind  interferes  with  the  execution  of  the 
analysis;  for  example,  with  the  operations  of  solution,  filtration,  and  evapora- 
tion. Alloys  do  not  contain  organic  matter  or  water;  and  therefore  the  closed- 
tube  test  need  not  be  applied  to  them. 

27 


28  PREPARATION  OF  THE  SOLUTION.  P.I 

2.  Blackening  accompanied  by  a  burnt  odor  or  by  the  formation  of  a  tarry 
deposit  shows  organic  matter.    Blackening  alone  does  not  show  it;  for  copper, 
cobalt,  and  nickel  salts  may  turn  black  on  heating,  owing  to  the  formation 
of  the  black  oxides. 

3.  It  is  usually  desirable  to  determin  whether  water  is  a  constituent  of  the 
substance,  and,  if  so,  whether  it  is  present  in  large  or  small  proportion.     This 
can  be  done  with  a  fair  degree  of  delicacy  by  the  closed-tube  test,  provided 
care  be  taken  to  keep  the  upper  part  of  the  tube  cool  during  the  first  of  the 
heating.     Water  may  be  present  as  so-called  water  of  constitution,  as  in 
FeOsHs  or  NagHPO^  as  water  of  crystallization,  as  in  MgSOi.THsO;  as  enclosed 
water,  as  in  some  hydrated  silicates  like  the  zeolites  or  as  mother-liquor  within 
crystals;  and  as  hygroscopic  moisture  on  the  surface.      Water  of  constitution 
may  be  expelled  only  at  a  fairly  high  temperature,  while  in  the  other  forms  it 
is  seldom  retained  above  200°. 

4.  The  closed-tube  test  may  also  furnish  evidence  of  the  presence  of  certain 
basic  and  acidic  constituents  when  they  are  present  in  considerable  quantity. 
Thus  all  ammonium  salts  and  mercury  compounds  are  volatilized  much  below 
a  red  heat.    Ammonium  salts  and  the  chlorides  of  mercury  giv  a  white  subli- 
mate.   Most  other  mercury  compounds  giv  a  gray  one,  consisting  of  minute 
globules  of  mercury,  made  visible  by  a  lens  or  by  rubbing  with  a  wire.    Metallic 
As,  As2O3,  and  As2Ss  are  also  readily  volatilized,  forming  black,  white,  and  yellow 
sublimates,  respectivly.     Of  the  acid-forming  elements  or  groups,  free  sulfur  or  a 
persulfide  is  shown  by  a  sublimate  of  reddish-brown  drops,  changing  to  a  yellow 
solid  on  cooling,  and  accompanied  by  odor  of  SO2J  a  moist  sulfide,  by  the  odor 
of  £[28;  a  nitrate  or  nitrite,  by  brown  vapors  of  NO2J  free  iodin  or  a  decom- 
posable iodide,  by  a  black  sublimate  of  Iz  and  by  its  violet  vapor;   a  sulfite, 
by  the  odor  of  SOz',   a  peroxide,  chlorate,  or  nitrate,  by  evolution  of  oxygen, 
recognized  by  its  inflaming  a  glowing  wood-splinter  held  in  the  tube;  and  a 
carbonate  or  oxalate,  by  the  evolution  of  CO2,  recognized  by  its  causing  tur- 
bidity in  a  drop  of  Ba(OH)2  solution. 

5.  If  the  substance  to  be  analyzed  is  a  liquid,  it  is  desirable  to   determin 
by  evaporation  how  much,  if  any,  solid  substance  is  present  in  it;  for  enough 
must  be  taken  for  analysis  to  enable  small  quantities  of  the  basic  constituents 
to  be  detected.     Moreover,  if  it  is  dissolved  in  a  volatil  organic  solvent  the 
latter  must  be  removed  by  evaporation. 


P.S 


PREPARATION  OF  THE  SOLUTION. 


29 


PREPARATION    OF    THE    SOLUTION. 


TABLE  I. — PREPARATION  OF  THE  SOLUTION  IN  THE  CASE  OP  NON-METALLIC 

SUBSTANCES. 


Heat  the  substance  with  water  and  dilute  HNOz  (P.  2). 


If  it  all 
dissolves, 
treat  the 


If  it  does  not  all  dissolve,  add  more  HNOz,  evaporate,  dry  completely, 
add  dilute  HN03  (P.  5). 


solution 
by  P.  11. 

Substances 
decom- 
posed: 
many  sul- 
fides  and 
silicates. 
Treat  the 
solution  by 
P.  11. 

Residue:*  a.  Sb2O6,  H2SnO3,  MnQj,  PbO2,  HgS,  PbCr04, 
BaCrO4.   b.  C,  A12O3,  Cr2O3,  AgCl,  CaF2,  PbS04,  BaSO4, 
SrSO4,  SiO2,  and  many  silicates  and  fluosilicates. 
Heat  with  HCl  or  HCl  and  HNOs,  evaporate,  add  dilute 
HCl  (P.  4). 

Solution: 
substances 
under  a. 
Treat  by 
P.  21. 

Residue:  substances  under  b. 
Heat  with  #2£04  and  HF,  evaporate  off 
the  HF,  add  water,  boil  (P.  5). 

Gas: 

SiF4. 

Residue: 
Pb,  Ba,  Sr,(Cr), 
as  sulfates. 
Treat  by  P.  6. 

Solution  : 
other  elements 
as  sulfates. 
Treat  by  P.  11. 

*  Only  the  more  common  substances  that  are  likely  to  be  present  in  the  residue  are  here  mentioned. 

Procedure  2. — Treatment  of  Non-Metallic  Substances  Free  from 
Organic  Matter. — Weigh  out  on  a  rough  balance  1  g.  of  the  finely 
powdered  substance  (see  Note  1),  add  to  it  in  a  casserole  20  cc. 
water,  heat  the  mixture  to  boiling  if  there  is  a  residue,  and  test  the 
solution  with  litmus  paper.  Add  to  the  mixture  6-normal  HN03,  a 
few  drops  at  a  time,  till  after  shaking  it  becomes  distinctly  acid. 
Note  whether  there  is  an  odor  or  effervescence.  Then  add,  without 
filtering  out  any  residue  and  without  further  heating,  just  5  cc. 
6-normal  HN03. 

If  the  substance  has  dissolved  completely,  treat  the  solution  by  P.  11. 

If  the  substance  has  not  dissolved  completely,  treat  the  mixture, 
without  filtering  out  the  residue,  by  P.  3.  (See  Note  8.) 

Notes. — 1.  In  order  that  difficultly  soluble  substances  may  be  dissolved, 
the  substance  must  be  reduced  to  a  very  fine  powder.  This  is  usually  best 
accomplished  by  grinding  the  substance,  a  small  quantity  at  a  time,  in  a  por- 
celain or  agate  mortar.  With  hard  substances,  and  in  general  with  minerals,  an 


30  PREPARATION  OF  THE  SOLUTION.  P.  2 

agate  mortar  should  be  used.    As  such  a  mortar  is  likely  to  be  broken  by  a 
blow,  the  substance  should  be  ground,  not  pounded,  in  it. 

2.  The  quantity  of    the  substance  taken  for  analysis  should  always  be 
approximately  known;    for  a  good  qualitativ  analysis  should  not  only  show 
the  presence  or  absence  of  the  various  elements  in  the  substance,  but  should 
enable  their  relativ  quantities  to  be  estimated.    Since  1  or  2  mg.  of  almost  any 
element  can  be  detected  by  this  system  of  analysis,  the  presence  of  0.1-0.2% 
of  an  element  will  be  detected  when  one  gram  of  substance  is  taken,  and  this 
degree  of  delicacy  is  ordinarily  sufficient.     If  much  more  than  this  quantity 
is  taken,  the  precipitates  may  be  so  large  that  much  time  is  consumed  in  fil- 
tering and  washing  them.     Moreover,  the  directions  given  for  some  of  the 
separations  are  based  on  the  assumption  that  not  more  than  500  mg.  of  any  one 
constituent  is  present. 

3.  When  the  substance  dissolves  only  partly  in  water,  it  is  not  worth  while 
to  filter  off  the  residue  and  analyze  it  and  the  solution  separately,  unless  special 
information  in  regard  to  the  soluble  constituents  is  desired.     It  is,  therefore, 
directed  to  treat  at  once  with  HNOs.     The  mixture  is  not  heated  after  addition  of 
the  acid,  so  as  to  avoid  oxidizing  mercurous,  arsenous,  and  ferrous  salts.     Only 
20  cc.  of  water  are  used  so  that  the  acid  may  be  strong  enough  to  prevent  the 
hydrolysis  of  salts  of  bismuth  and  tin,  and  thus  ensure  their  solution. 

4.  Just  5  cc.  6-normal  HNOs  must  be  present  in  order  that  the  acid  con- 
centration may  be  properly  adjusted  in  the  subsequent  H2S  precipitation.     For 
this  reason,  when  the  solution  is  alkaline  or  when  a  substance  (like  an  undissolved 
oxide  or  carbonate)  which  neutralizes  the  acid  is  present,  the  solution  is  made 
distinctly  acid  before  adding  the  5  cc.  of  HNOs. 

5.  If  the  aqueous  solution  has  an  alkaline  reaction,  the    addition  of  an 
acid  may  cause  precipitation  of  any  substance  held  in  solution  by  an  alkaline 
solvent;  for  example,  sulfur  or  sulfides  of  the  tin-group  from  an  alkaline  sulfide 
solution;  silver  chloride  or  cyanide  from  a  potassium  cyanide  solution;  silicic 
acid  from  sodium  silicate  solution;    or  basic  hydroxides  from  solutions  in 
alkalies.    These  last  substances  redissolve  when  the  excess  of  HNOs  is  added. 

6.  An  acid  reaction  of  the  aqueous  solution  towards  litmus  is  due  to  hy- 
drogen-ion, which  may  arise  from  free  acid,  from  an  acid  salt  of  a  strong  acid, 
or  (by  hydrolysis)  from  a  neutral  salt  of  a  strong  acid  and  a  weak  base.    An 
alkaline  reaction  is  due  to  hydroxide-ion,  which  may  arise  from  a  soluble  hy- 
droxide, or  (by  hydrolysis)  from  a  carbonate,  sulfide,  phosphate,  borate,  cyanide, 
or  a  salt  of  some  other  weak  acid. 

7.  When  the  acid  is  added  to  the  aqueous  solution,  the  evolution  of  any 
gas  and  its  odor  should  be  noted,  since  this  indicates  the  nature  of  the  acidic 
constituents  present.      Thus   carbonates  effervesce  with   evolution  of  CO2,* 
sulfides  produce  the  odor  of  H2S;  sulfites  and  thiosulfates,  that  of  S02;  and 
cyanides,  that  of  HCN. 

8.  There  are  certain  compounds  (especially  those  of  antimony  and  tin,  the 
oxides  of  iron,  MnO2,  and  BaCrO4)  which  dissolve  more  rapidly  or  in  larger 
quantity  in  HC1  than  in  HNO3.    When  a  substance  fails  to  dissolve  completely 
in  the  dilute  HNO8  and  seems  likely  from  its  appearance  or  behavior  to  contain 
any  of  these  compounds,  it  is  therefore  well  to  attempt  to  prepare  an  HC1  solu- 
tion of  it  by  proceeding  as  follows:    To  a  fresh  1  g.  sample  of  the  finely  pow- 
dered substance  in  a  small  flask  add  just  5  cc.  6-normal  HC1;  heat  the  mixture 


P.  t  PREPARATION  OF  THE  SOLUTION.  31 

nearly  to  boiling  for  3-4  minutes,  covering  the  flask  with  a  watch-glass  to  pre- 
vent evaporation;  and  then  add  10  cc.  water.  If  the  substance  has  dissolved 
completely,  reject  the  HNO3  solution  and  residue,  and  treat  the  HC1  solution 
by  P.  21.  Otherwise,  reject  the  HC1  solution  and  residue,  and  treat  the  HNO8 
solution  and  residue  by  P.  3. — As  to  the  reasons  for  recommending  the  use  of 
HNO3,  rather  than  of  HC1,  as  the  usual  procedure,  see  Note  4,  P.  3. 

9.  The  following  general  statements  may  be  made  in  regard  to  the  solubility 
of  substances  in  water  and  dilute  acids: 

All  the  ordinary  salts  of  sodium,  potassium,  and  ammonium  are  readily  soluble 
in  water. 

The  salts  of  mercurous  and  mercuric  mercury,  bismuth,  antimony,  and  tin 
are  hydrolyzed  by  water  with  precipitation  of  basic  salts,  which  dissolve  readily 
in  dilute  HNO3  or  HC1. 

The  nitrates,  nitrites,  chlorates,  and  acetates  of  all  the  elements  are  readily 
soluble  in  water  (except  certain  basic  nitrates  and  acetates). 

The  hydroxides,  carbonates,  phosphates,  borates,  arsenates,  and  arsenites  of 
all  the  elements  except  the  alkalies  are  only  slightly  soluble  in  water,  but  dis- 
solve readily  in  dilute  HNO3  or  HC1.  (Ba(OH)2,  Sr(OH)2,  and  Ca(OH)2  are, 
however,  fairly  soluble  in  water.) 

The  chlorides,  bromides,  iodides,  and  thiocyanates  of  all  the  elements  except 
lead,  silver,  and  mercury,  and  the  sulfates  of  all  the  elements  except  calcium, 
strontium,  barium,  lead,  mercury,  and  silver  are  readily  soluble  in  water.  (In 
regard  to  the  solubility  of  these  and  other  salts  of  the  alkaline  earths  and  of 
silver  and  lead,  see  the  Table  of  Solubilities  of  Slightly  Soluble  Substances  on 
page  124.) 

Procedure  3. — Treatment  of  Non-Metallic  Substances  not  dissolved 
by  Dilute  Nitric  Add  and  of  Alloys. — If  the  substance  is  non-metallic 
and  has  not  dissolved  in  dilute  HN03,  to  the  mixture  obtained  in 
P.  2  add  5  cc.  16-normal  HN03,  and  evaporate  just  to  dryness. 

If  the  substance  is  an  alloy,  convert  it  into  a  form  offering  a  large 
surface  and  treat  0.5  g.  of  it  in  a  casserole  with  10  cc.  6-normal 
HN03.  Cover  the  dish  with  a  watch-glass,  heat  the  mixture  nearly 
to  boiling  as  long  as  any  action  continues,  adding  a  little  16-normal 
HNOi  if  action  is  renewed  thereby,  or  a  little  water  if  crystalline  salts 
have  separated,  and  then  evaporate  just  to  dryness. 

Heat  the  residue  obtained  in  either  case  at  100-130°  until  it  is 
perfectly  dry,  by  keeping  the  casserole  in  motion  over  a  small  flame. 
Loosen  the  residue  from  the  dish  and  rub  it  to  a  fine  powder  with 
a  pestle;  add  to  it  just  5  cc.  6-normal  HN03,  cover  the  dish,  and 
warm  the  mixture,  taking  care  that  none  of  the  acid  evaporates. 
Dilute  with  20  cc.  water,  heat  to  boiling,  and  note  whether  there  is 
any  residue  (see  Note  7).  If  there  is  a  residue,  filter  it  out  and  wash 
it.  (Residue,  P.  4;  solution,  P.  11.) 

Notes. — 1.  On  heating  the  HNOs  solution,  the  presence  of  sulfides  is  indi- 
cated by  the  separation  of  sulfur  as  a  spongy  or  pasty  mass,  which  floats  on 


32  PREPARATION  OF  THE  SOLUTION.  P.  3 

the  surface  and  may  be  removed  by  means  of  a  spatula  or  rod;  and  the  presence 
of  iodides  is  shown  by  the  liberation  of  free  iodin,  which  may  separate  as  a 
black  precipitate,  which  imparts  a  brown  color  to  the  solution,  and  which 
gives  rise  to  violet  vapors  above  it. 

2.  When  a  silicate  is  decomposed  by  acid,  silicic  acid  may  separate  as  a 
gelatinous  precipitate,  but  even  then  a  part  of  it  always  remains  in  solution, 
mainly  as  a  colloid.    When  thoroly  dried  at  100-130°,  it  is  partially  dehydrated 
and  becomes  entirely  insoluble.    The  HNOs  acid  solution  is  therefore  evapo- 
rated to  dryness  and  the  residue  is  heated  at  100-130°,  in  order  to  remove  the 
silica  at  this  point;  for,  if  it  were  not  removed,  it  would  appear  as  a  gelatinous 
precipitate  at  some  later  stage  of  the  analysis;    thus,  if  it  did  not  separate 
earlier,  it  would  be  precipitated  by  NH4OH  together  with  the  iron  group  and 
might  then  be  mistaken  for  aluminum  hydroxide.    In  the  case  of  nonmetallic 
substances  which  cannot  contain  silica,  the  heating  may  be  omitted. 

3.  If  the  substance  is  nonmetallic,  the  residue  insoluble  in  HNOs  probably 
consists  of  one  or  more  of  the  substances  whose  formulas  are  given  in  Table  I 
under  a  and  6.     Other  less  common  insoluble  substances  are  anhydrous  chro- 
mium salts,  phosphate  of  tin,  ferrocyanide  of  iron,  and  silicon  carbide. 

4.  In  dissolving  nonmetallic  substances  HC1  may  be  used  in  place  of  HNOs. 
Each  of  these  acids  has  advantages  and  disadvantages  of  its  own,  as  follows: 
HNO3  dissolves,  owing  to  its  oxidizing  power,  many  sulfides  not  attacked  by 
HC1,  but  fails  to  dissolve  certain  substances,  especially  MnO2,  Sb2O5,  ^SnOs,  and 
BaCrO4,  which  dissolve  in  HC1.     HC1  may  cause  the  precipitation  of  chlorides 
of  the  silver  group;   while  strong  HNOs  on  heating  oxidizes  sulfides  partially 
to  sulfates,  and  may  cause  the  precipitation  of  lead,  barium,  strontium,  and 
calcium  sulfates;  thus  in  either  case  making  it  sometimes  impossible  to  deter- 
min  whether  complete  decomposition  has  resulted.    HNOs  oxidizes  mercurous, 
arsenous,  antimonous,  stannous,  and  ferrous  compounds  to  the  higher  state 
of  oxidation;    consequently  almost  all  the  antimony  and  tin  will  usually  be 
found  in  the  residue  insoluble  in  dilute  HNOs  after  evaporation,  all  the  mercury 
will  be  in  the  H2S  precipitate,  and  sulfur  will  always  be  precipitated  by  H2S 
when  iron  is  present.     When  HC1  is  used  as  a  solvent,  mercury  and  arsenic 
in  the  arsenous  form  would  be  wholly  or  partly  lost,  owing  to  the  volatility 
of  their  chlorides,  in  the  subsequent  evaporation,  which  is  necessary  in  order 
to  remove  silica.    For  this  last  reason,  and  for  the  reason  that  the  procedure 
is  a  more  genera1  one  in  that  it  provides  for  the  solution  of  alloys  and  of  a 
larger  proportion  of  nonmetallic  substances  and  for  the  isolation  of  the  silver 
group,  the  use  of  HNOs  is  here  recommended. 

5.  Alloys  can  not  ordinarily  be  powdered  by  grinding  in  a  porcelain  or 
agate  mortar.    They  may  usually  be  converted  into  a  form  that  offers  a  large 
surface  by  hammering  in  a  steel  mortar,  filing  with  fine  steel  file,  shaving  with 
a  knife,  or  converting  into  turnings  with  a  lathe.     Only  0.5  g.  of  an  alloy  is 
taken  for  analysis;   for,  owing  to  the  absence  of  acidic  constituents,  the  same 
quantity  of  basic  elements  is  contained  in  a  smaller  amount  of  substance. 

6.  By  the  treatment  of  alloys  with  strong  HNOs,  all  the  more  common 
elements  are  dissolved  by  strong  HNOs  except  antimony,  tin,  and  silicon. 
These  are  oxidized  to  antimonic acid  (Sb2O5.nH20),metastannic  acid  (n^SnOg), 
and  silicic  acid  (HzSiOs),  which  separate  at  once  as  white  amorphous  pre- 
cipitates when  considerable  amounts  of  these  elements  are  present.    Certain 


P.  3  PREPARATION  OF   THE  SOLUTION.  33 

nitrates,  especially  that  of  lead,  may  separate  in  crystalline  form  from  the 
strong  HNOs,  but  these  dissolve  upon  adding  water  and  heating  to  boiling. 

7.  In  the  case  of  an  alloy  the  evaporation  to  dryness  and  heating  at  100- 
130°  serve  to  partially  dehydrate  the  hydroxides  of  silicon,  tin,  and  antimony, 
whereby  they  are  rendered  nearly  insoluble  in  HNOs.     This  makes  possible 
a  conclusion  in  regard  to  their  presence  or  absence.     Thus,  if  after  having 
thoroly  dried  the  mixture  at  this  temperature  there  is  no  residue  insoluble 
in  the  HNOs,  it  shows  the  absence  of  silicon  and  tin  in  quantity  as  large  as 
1  mg.,  and  that  of  antimony  in  quantity  as  large  as  2  or  3  mg.    The  fact  must 
not  be  overlooked,  however,  that  in  the  dehydrated  form  even  a  very  small 
residue  or  slight  turbidity  may  correspond  to  an  appreciable  quantity  of  one 
of  these  elements.    Therefore,  if  no  residue  can  be  seen,  rub  the  sides  of  the 
dish  gently  with  the  rubber-covered  end  of  a  glass  rod,  pour  into  a  small  flask, 
allow  the  liquid  to  stand  2  or  3  minutes,  and  note  whether  there  is  any  residue 
whatever.    The  knowledge  that  tin  is  absent  enables  the  subsequent  procedures 
for  the  detection  of  this  element  to  be  omitted.    The  subsequent  procedures 
for  antimony  may,  in  the  absence  of  a  residue,  also  be  omitted,  provided 
quantities  as  small  as  3  mg.  are  not  to  be  tested  for.    In  addition  to  the  hy- 
droxides named  above,  the  residue  may  also  contain  a  considerable  quantity 
of  stannic  phosphate  or  arsenate  when  tin  and  pho  phorus  or  arsenic  are  simul- 
taneously present,  or  of  bismuth  hydroxide  when  both  antimony  and  bismuth 
are  present;    also  small  quantities  of  various  other  elements  enclosed  in  a 
residue  consisting  of  the  substances  already  mentioned. 

8.  The  hydroxides  of  antimony,  tin,  and  silicon  usually  separate  also  in 
the  treatment  of  nonmetallic  substances  with  HNOs  when  the  corresponding 
elements  are  present;   but  the  nonexistence  of  a  residue  must  not,  except  in 
the  case  of  silicon,  be  regarded  as  conclusiv  evidence  of  their  absence  in  such 
substances.     For  the  presence  of  certain  acidic  constituents,  such  as  chloride 
or  sulfate,  may  cause  a  considerable  quantity  of  tin  or  antimony  to  dissolve. 

9.  A  black  or  metallic  residue  insoluble  in  HNOs,  obtained  in  the  case  of 
an  alloy,  may  contain  carbon  or  carbides,  certain  alloys  of  iron,  such  as  ferro- 
chrome  or  ferrosilicon,  gold,  or  any  of  the  platinum  metals.     If  there  is  no 
such  residue,  it  shows  the  absence  of  gold  and  platinum. 

Procedure  4. — Treatment  of  the  Residue  Insoluble  in  Nitric  Acid.— 
To  the  residue  insoluble  in  HNOs  (P.  3)  in  a  casserole  add  gradually 
5  cc.  12-normal  HC1,  and  heat  as  long  as  action  continues,  adding 
more  acid  if  necessary.  If  the  substance  does  not  dissolve  completely 
in  HC1,  add  to  the  mixture  without  filtering  a  few  drops  of  16-normal 
HN03,  and  heat  gently  as  long  as  action  continues,  adding  more  of  the 
acids  if  necessary. 

Evaporate  this  solution  in  HC1  alone,  or  in  HC1  and  HN03,  without 
filtering  off  any  residue,  just  to  dryness,  taking  care  not  to  overheat 
the  dry  residue.  Add  to  the  residue  5  cc.  6-normal  HC1,  measured 
in  a  small  graduate,  and  10  cc.  water;  boil  gently  for  a  few  minutes 
if  there  is  a  residue:  filter,  wash  the  residue  thoroly  with  boiling 
water,  and  treat  it  by  P.  5  (or  P.  7).  Unite  the  filtrate  with  the 


34  PREPARATION  OF  THE  SOLUTION.  P.  4 

HNOs  solution  obtained  in  P.  3  after  treating  the  latter  by  P.  11; 
neutralize  half  the  acid  in  the  mixture  by  adding  just  5  cc.  6-normal 
NH40H;  and,  without  filtering  off  any  precipitate,  treat  the  mixture 
by  P.  21. 

Notes. — 1.  Of  the  substances  that  may  be  present  in  the  residue  undis- 
solved  by  HNOs  (see  Table  I),  the  peroxides  of  manganese  and  lead  and  the 
chromates  of  barium  and  lead  are  reduced  and  dissolved  by  concentrated  HC1. 
Antimonic  acid  (HSbOs)  and  stannic  acid  (H^SnOs)  are  also  dissolved  by  it. 
Upon  the  addition  of  HNOs,  whereby  the  strongly  oxidizing  mixture  known  as 
aqua  regia  is  produced,  gold,  platinum,  and  mercuric  sulfide  are  entirely  dissolved; 
and  silver  compounds,  such  as  AgBr,  Agl,  and  AgCN,  are  converted  into  AgCl. 
The  chloride  of  silver  and  the  sulfates  of  strontium  and  lead  dissolve  in  large 
quantity  in  the  concentrated  acids,  but  only  in  much  smaller  quantity  in  the 
small  amount  of  dilute  HC1  added  after  the  evaporation.  Some  of  the  other 
substances  that  may  be  in  the  residue,  especially  the  oxides  and  certain  sili- 
cates, are  slowly  attacked  by  the  strong  acids,  but  the  solvent  action  is  not 
rapid  enough  to  make  this  a  practicable  method  of  getting  them  into  solution. 

2.  The  solution  is  evaporated  to  remove  the  large  quantity  of  acid  which 
would  otherwise  interfere  with  the  £[28  precipitation.    Care  is  taken  not  to  over- 
heat the  dry  residue,  so  as  to  avoid  loss  of  mercury,  antimony,  and  tin  by 
volatilization  of  their  chlorides.    A  measured  quantity  of  HC1  is  then  added; 
and,  after  mixing  this  HC1  solution  with  the  HNOs  solution,  an  equivalent 
quantity  of  NB^OH  is  added,  in  order  to  produce  the  acid  concentration  required 
for  the  H2S  precipitation.     Only  10  cc.  of  water  are  added  to  the  HC1  solution 
at  first,  so  as  to  prevent  the  precipitation  of  SbOCl  and  thus  make  it  possible 
to  determin  whether  the  substance  has  been  completely  decomposed. 

3.  If  the  original  substance  was  an  alloy,  a  residue  after  the  treatment 
with  HC1  and  HNOs  is  likely  to  consist  of   metastannic  or  silicic  acid    or  of 
carbon,  a  platinum  metal,  or  an  alloy  of  iron  with  chromium,  silicon,  etc.    It 
is  best  treated  with  H2S04  and  HF  by  P.  5,  in  order  to  test  for  and  remove 
silica  and  to  dissolve  metastannic  acid  and  iron-alloys.     If  a  black  or  metallic 
residue  still  remains,  it  may  be  tested  for  graphite  by  rubbing  a  dried  portion 
on  the  fingers  or  on  paper;   and  to  bring  it  into  solution  the  remainder  may 
then  be  fused  with  Na2O2  in  a  nickel  crucible,  the  mass  treated  with  water 
and  HC1,  and  the  solution  analyzed  as  usual,  except  that  nickel  cannot  be 
tested  for. 

4.  If  the  original  substance  was  an  alloy  and  a  large,  nonmetallic  residue 
remains  after  treatment  with  HNOs  (P.  3),  it  is  sometimes  advantageous, 
instead  of  treating  it  by  P.  4,  to  analyze  the  residue  separately  by  the  following 
procedure,  by  which  a  large  quantity  of  metastannic  acid  is  more  readily  dis- 
solved:  Add  to  the  residue  in  a  casserole  3-4  cc.  of  96  %  H2SO4  and  heat 
under  the  hood  until  the  acid  has  evaporated  to  a  volume  of  about  2  cc.    Cool, 
add  an  equal  volume  of  water,  cool  again,  add  5  cc.  HC1  to  dissolve  antimonic 
oxide,  and  heat  to  boiling.    Cool  completely,  filter  if  there  is  a  residue  (which 
may  consist  of  silicic  acid),  and  add  the  acid  solution  drop  by  drop,  with  con- 
stant shaking,  to  a  mixture  of  10  cc.  ammonium  monosulfide,  1  cc.  ammonium 
polysulfide,  and  10  cc.  15-normal  NH4OH  in  a  flask.     Cover  the  flask  and 
digest  for  a  few  minutes  on  a  steam  bath.    Filter  out  the  precipitate,  which 


P.  5  PREPARATION  OF  THE  SOLUTION.  35 

may  consist  of  small  quantities  of  sulfides  of  the  copper  and  iron  groups.  Dilute 
the  filtrate,  and  make  it  slightly  acid  with  HC1.  Shake  to  coagulate  the  pre- 
cipitate, filter,  and  wash  with  hot  water.  Analyze  the  precipitate  for  the  tin- 
group  by  P.  41;  reject  the  filtrate  or  test  it  for  phosphate  by  P.  104. 

*Procedure  5. — Fluoride  Treatment  of  the  Residue  Insoluble  in  the 
Common  Acids. — Transfer  to  a  platinum  crucible  (see  Notes  1  and  2) 
the  residue  after  treatment  with  acids  (P.  4),  add  2  cc.  of  96%  H2S04 
from  a  graduate,  heat  with  a  moving  flame  until  white  fumes  are 
given  off,  and  cool  completely. 

To  test  for  silicate,  add  carefully  from  a  lead  or  hard-rubber  tube 
capped  with  a  rubber  nipple  pure  48%  HF  drop  by  drop  until  5-6 
drops  have  been  added,  and  warm  the  mixture  over  a  steam  bath. 
(Formation  of  gas  bubbles,  presence  of  SILICA  or  SILICATE.) 

Then  add  2-5  cc.  more  pure  48%  HF,  cover  the  crucible  with  a 
platinum  cover,  digest  on  a  steam  bath  for  about  15  minutes  unless 
the  residue  dissolves  more  quickly;  remove  the  cover,  and  evaporate 
under  a  hood  until  dense  white  fumes  of  H2S04  are  given  off,  taking 
care  to  avoid  spattering.  (See  Note.  3.)  [Unless  it  is  known  from  the 
presence  of  solid  substance  at  this  point  or  from  other  indications  that 
the  residue  treated  with  H2S04  and  HF  contained  other  constituents 
than  silica,  determin  this  by  evaporating  off  the  H2S04  under  a  hood, 
taking  care  not  to  ignite  the  dry  residue.  If  a  significant  residue 
remains,  add  from  a  graduate  1.5  cc.  of  96%  H2S04,  and  heat  until  the 
residue  is  redissolved,  not  allowing  the  acid  to  evaporate.]  Cool,  pour 
the  contents  of  the  crucible  into  10  cc.  water,  and  rinse  out  the  contents 
with  a  little  water.  Boil  to  dissolve  slowly  dissolving  sulfates;  cool, 
shake,  filter,  and  wash  the  residue,  first  with  6-normal  H2S04  and  then 
with  a  little  water.  (Residue,  P.  6;  filtrate,  P.  11.) 

Notes. — 1.  A  student  using  this  procedure  for  the  first  time  should  work  under 
the  direct  supervision  of  an  instructor.  Great  care  must  be  taken  not  to  breathe 
the  fumes  of  HF  nor  to  get  it  on  the  hands;  for  it  is  extremely  irritating  and  pro- 
duces dangerous  burns. 

2.  Whenever  a  residue  or  precipitate  has  to  be  transferred  from  a  filter  to  a  cru- 
cible in  which  it  is  to  be  ignited  or  fused,  it  is  best  to  separate  it  as  far  as  possible 
from  the  filter-paper,  and  then  to  incinerate  the  part  of  the  paper  to  which  much 
residue  adheres  by  rolling  it  up,  winding  a  platinum  wire  around  it  in  the  form  of  a 
spiral,  and  heating  it  in  a  gas-flame  till  the  carbon  is  all  burnt  off. 

3.  When  a  liquid  is  to  be  evaporated  in  a  crucible,  it  is  well  to  heat  it  within 
a  larger  iron  crucible,  which  serves  as  an  air-bath.     The  smaller  crucible  may  be 
supported  upon  a  nicrome  triangle  set  into  holes  bored  in  the  side  of  the  iron  crucible, 
or  upon  a  circular  disk  of  asbestos-board  with  a  round  hole  cut  out  in  the  middle  and 
slots  cut  out  along  the  sides. 

*  If  the  use  of  a  platinum  crucible  or  of  hydrofluoric  acid  is  impracticable,  the  lees  satisfactory, 
alternativ  method  described  in  P.  7  may  be  employed  (see  Note  8,  P.  5). 


36  PREPARATION  OF  THE  SOLUTION.  P.  6 

4.  The  test  for  silica  or  silicate  depends  on  the  formation  of  SiF4  gas,  which 
is  insoluble  in  strong  H2S04,  but  dissolves  in  water  in  the  presence  of  HF  with 
formation  of  fluosilicic  acid,  H2SiFe.     With  free  silica   the  evolution  of  gas 
takes  place  in  the  cold;  but  with  slowly  decomposing  silicates,  such  as  feldspar, 
the  test  is  obtained  only  upon  warming.     A  few  silicates  are  not  acted  upon 
by  HF  and  H2SO4,  and,  of  course,  do  not  show  the  test  for  silica  at  this  point. 
The  test  is  delicate  enough  to  enable  1  mg.  of  silica,  whether  free  or  in  a  de- 
composable silicate,  to  be  detected.     Moreover,  after  the  substance  has  been 
treated  with  acids  as  in  P.  4  and  warmed  with  H2SO<,  an  evolution  of  gas  with 
HF  is  not  produced  with  the  compounds  of  any  element  other  than  silicon. 
It  should  be  borne  in  mind  that  a  small  quantity  of  silica  will  be  introduced 
if  ordinary  niters  (which  have  not  been  washed  with  HF;  have  been  employed 
and  have  been  destroyed  by  acids  or  by  ignition,  or  if  a  strongly  alkaline  solu- 
tion has  been  boiled  in  glass  vessels. 

5.  Since  glass  and  porcelain  consist  of  silicates  which  are  readily  attacked 
by  HF,  this  acid  must  not  be  allowed  to  come  into  contact  with  these  materials. 
In  handling  cold  HF  solutions,  vessels  and  funnels  of  celluloid  or  paraffin  or 
of  glass  coated  with  paraffin  may  be  used;   but  platinum  vessels  must  be  em- 
ployed when  the  solutions  are  to  be  heated.    Care  must  be  taken  not  to  intro- 
duce into  a  platinum  vessel  any  solution  containing  chlorin  or  bromin  or  any 
acid  mixture  containing  nitrates  and  chlorides  by  which  chlorin  would  be 
evolved.     Platinum  is  so  slowly  attacked  by  hot  concentrated  H2SO4  that 
even  when  2-3  cc.  of  the  acid  are  rapidly  evaporated  in  a  crucible  less  than 
0.5  mg.  passes  into  solution. 

6.  The  digestion  with  HF  decomposes  most  silicates  and  dissolves    silica. 
The  subsequent  evaporation  with  H2S(>4  expels  the  excess  of  HF  and  decom- 
poses the  fluorides  produced,  as  well  as  some  other  substances  that  may  have 
been  left  undissolved  by  the  HNO3  and  HC1.    The  H2SO4  solution  is  diluted 
with  a  small  quantity  of  water  so  as  to  cause  the  complete  precipitation  of 
BaSO4,  SrSC>4,  and  PbSO4.      These  sulfates  are  moderately  soluble  in  strong 
H2SC>4  and  may  not  appear  till  after  dilution.     The  addition  of  much  water 
is  avoided,  since  SrSC>4  and  PbSC>4  are  somewhat  soluble  in  water;    and  the 
residue  is  washed  with  dilute  ^864  for  the  same  reason.    The  solution  is 
boiled  so  as  to  dissolve  anhydrous  sulfates,  such  as  those  of  aluminum  and 
iron. 

7.  The  residue  insoluble  in  dilute  H2SC>4  contains  as  sulfates  all  the  barium, 
strontium,  and  lead,  and  all  of  the  calcium  in  excess  of  5-10  mg.,  left  undis- 
solved by  HNOa  and  HC1;   more  or  less  of  the  chromium  (according  as  the 
H2SO4  has  been  more  or  less  strongly  heated)  as  a  pink  anhydrous  sulfate; 
and  part  of  the  bismuth  as  basic  sulfate  and  antimony  as  antimonic  hydroxide, 
when  much  of  these  elements  was  left  undissolved  by  the  previous  treatments 
with  acids.     The  residue  may  also  contain  still  undecomposed  substances, 
especially  the  following:  silver  chloride;  corundum,  Al2Os;  chromite,  FeCr2O4j 
cassiterite,  SnO2;    some  anhydrous  silicates  and  fluosilicates,  such  as  cyanite 
or  andalusite    (Al2SiOs)  and  tourmalin;   graphite  and  carbides;  and  certain 
compounds  of  the  rarer  elements.     For  a  method  of  bringing  some  of  these 
substances  into  solution,  see  Note  3,  P.  7. 

8.  If  the  use  of  a  platinum  crucible  or  of  hydrofluoric  acid  is  impracticable, 
the  residue  insoluble  in  HC1  and  HNOs  may  be  fused  in  a  nickel  crucible  with 


P.  6  PREPARATION  OF  THE  SOLUTION.  37 


,  as  described  in  P.  7,  instead  of  being  treated  by  P.  5-6.  This  is, 
however,  a  far  less  satisfactory  method  of  analysis  for  the  following  reasons. 
Compounds  of  the  alkali  elements  are  used  as  a  flux;  nickel  is  introduced 
from  the  crucible;  and  mercury  compounds  are  volatilized;  —  so  that  these 
elements  can  not  be  tested  for  in  the  subsequent  analysis.  Moreover,  the 
treatment  with  HF  and  H2SO4  is  almost  always  a  shorter  process,  since  when 
the  residue  consists  only  of  silica,  as  is  often  the  case  with  minerals,  no  further 
treatment  is  necessary,  and  since  in  other  cases  there  is  often  no  residue  to  be 
boiled  with  Na2COs  solution  (P.  6).  A  fusion  in  a  platinum  crucible  with 
Na2COa  would  be  less  objectionable;  but  this  is  not  possible,  unless  elements 
reducible  to  the  metallic  state  are  known  to  be  absent  in  the  residue  (see  Note 
5,  P.  7). 

Procedure  6.  —  Treatment  of  the  Residue  from  the  Fluoride  Treat- 
ment. —  Transfer  the  residue  insoluble  in  dilute  H2S04  (P.  5)  to  a 
casserole,  add  about  25  cc.  3-normal  Na2COs  solution,  cover  the 
casserole,  and  boil  gently  for  10  minutes.  Filter  and  wash  the  resi- 
due thoroly.  (Filtrate,  reject.)  Heat  the  residue  with  just  5  cc. 
HNOs  and  10-20  cc.  water.  Filter  out  any  undissolved  residue,  and 
treat  the  solution  by  P.  11,  subsequently  testing  it  only  for  lead, 
bismuth,  chromium,  barium,  strontium,  and  calcium. 

Notes.  —  1.  The  boiling  with  NasCOs  converts  into  carbonates  the  sulfates 
of  lead,  calcium,  strontium,  and  bismuth  completely,  and  at  least  80%  of 
the  sulfate  of  barium,  even  when  large  quantities  of  them  are  present.  A 
second  treatment,  which  should  be  applied  to  the  residue  if  there  are  indica- 
tions that  barium  is  present,  completely  decomposes  BaSO4.  The  carbonates 
dissolve  readily  in  HNOa.  Anhydrous  chromic  sulfate,  which  is  left  undis- 
solved by  dilute  H2S04  (P.  5)  as  a  fine  pink  or  gray  powder,  is  slowly  changed 
by  boiling  with  Na2COs  to  a  greenish  blue  hydroxide  which  dissolves  in  the 
HNOa,  leaving  behind  the  still  undecomposed  sulfate.  Antimonic  oxide  dis- 
solves only  to  a  small  extent  (2-4  mg.)  in  the  N^COs  solution  or  in  the  dilute 
HN03. 

2.  Any  residue  insoluble  in  HNOs  can  therefore  consist  only  of  barium 
or  chromic  sulfate,  of  antimonic  oxide,  or  of  some  of  the  original  substance 
still  undecomposed,  which  is  likely  to  consist  of  one  of  the  nativ  oxides  or 
silicates  mentioned  in  P.  5,  Note  5.  If  such  a  residue  is  obtained,  it  can  ordi- 
narily be  rendered  soluble  by  fusion  with  Na2COs,  K^COs,  and  KNOs,  as  de- 
scribed in  P.  7;  but  in  this  case  a  platinum  crucible  may  be  used  for  the  fusion, 
provided  the  residue  be  first  heated  with  HC1  to  extract  any  Sb20s  that  may 
be  present  and  provided  silver  is  not  found  present  in  the  H2S04  solution 
obtained  in  P.  5. 

Procedure  7.  —  Alternativ  Treatment  of  the  Residue  Insoluble  in 
the  Common  Acids.  —  If  the  use  of  HF  (P.  5)  is  impracticable,  transfer 
the  residue  insoluble  in  acids  (P.  4)  to  a  nickel  crucible  (see  Note  2, 
P.  5),  heat  until  the  residue  is  dry,  mix  the  residue  with  ten  to  twenty 
times  its  weight  of  anhydrous  Na2COs,  cover  the  crucible,  heat  strongly 


38  PREPARATION  OF  THE  SOLUTION.  P.  7 

over  a  powerful  burner,  preferably  within  a  cylinder  of  asbestos-paper, 
so  that  complete  fusion  takes  place,  and  continue  the  heating  for  10-20 
minutes.  If  dark  particles  of  undecomposed  substance  can  still  be 
seen,  add  gradually  in  small  portions  0.1-0.5  g.  of  solid  KN03,  and 
heat  strongly  for  several  minutes.  Cool,  boil  the  crucible  and  its 
contents  with  water  until  the  fused  mass  is  disintegrated,  filter,  and 
wash  the  residue  thoroly.  Warm  the  residue  with  HNO3  until  action 
ceases,  and  filter  out  any  still  undecomposed  substance. 

Test  one  half  of  the  carbonate  solution  for  acidic  constituents  as 
described  in  the  third  and  fourth  paragraphs  of  P.  131. 

Mix  a  tenth  of  the  HN03  solution  with  a  tenth  of  the  carbonate 
solution,  making  the  mixture  acid  with  HN03,  if  it  is  not  already 
so.  If  no  precipitate  forms,  mix  the  remainder  of  the  acid  solution 
with  the  remainder  of  the  carbonate  solution.  Add  3-5  cc.  HC1  (or 
more  if  the  solution  is  still  alkaline),  and  filter.  Test  the  precipitate 
for  lead  and  silver  by  P.  13-14.  Evaporate  the  solution,  and  heat  the 
residue  until  it  is  thoroly  dry  at  120-130°  by  keeping  it  in  motion  over 
a  small  flame.  Moisten  the  residue  with  16-normal  HN03,  and 
heat  again  till  the  residue  becomes  perfectly  dry.  Add  from  a 
graduate  just  5  cc.  6-normal  HNOs  and  about  20  cc.  water,  and 
heat  to  boiling.  Filter  out  any  residue,  dilute  the  filtrate  to  100  cc., 
and  treat  it  by  P.  21. 

If  a  precipitate  forms  on  mixing  the  small  portions  of  the  HNOs 
solution  and  the  carbonate  solution,  treat  these  solutions  separately 
as  described  in  the  preceding  paragraph,  uniting  the  precipitates 
formed  by  the  same  group-reagent  in  the  subsequent  analysis. 

Notes.  —  1.  Upon  fusion  with  sodium  carbonate  most  compounds  undergo 
metathesis,  the  acidic  constituent  of  the  compound  combining  with  the  sodium, 
and  the  basic  element  with  the  carbonate.  The  carbonate  formed  is,  however, 
sometimes  decomposed  by  heat  with  production  of  the  oxide  or  of  the  metal 
itself.  Acid-forming  oxides,  such  as  SiC>2,  As2O5,  and  less  rapidly  A^Oa,  expel 
CQj  from  the  carbonate  and  form  sodium  salts.  Such  reactions  are  illustrated 
by  the  following  equations: 

BaSO4  +Na2CO3  =Na2SO4  +BaCO3. 


4AgCl 
Si02 

After  the  treatment  with  water,  the  acidic  constituent  of  the  substance  is 
therefore  found  with  the  excess  of  carbonate  in  the  aqueous  extract,  while 
the  basic  element  remains  undissolved  by  the  water  and  passes  into  the  acid 
solu  ion.  The  first  and  third  reactions  are  examples  of  cases  where  the  aqueous 
and  acid  solutions  must  not  be  mixed,  for  upon  mixing  BaSC>4  or  AgCl  would 
again  be  formed. 


P.  7  PREPARATION  OF  THE  SOLUTION.  39 

2.  Of  the  basic  elements  that  may  be  present,  all  or  a  part  of  the  arsenic, 
antimony,  tin,  aluminum,  chromium    and  manganese  are  contained  in  the 
carbonate  solution;    and  this  solution  must  therefore  be  analyzed  for  basic 
elements.     Since  this  solution  may  also  contain  NagSiOs,  the  solution  after 
the  addition  of  acid  is  evaporated  to  dryness,  and  the  residue  is  heated    at 
100-130°,  in  order  to  dehydrate  the  silicic  acid  and  render  it  insoluble.    The  resi- 
due insoluble  in  HNOs  after  this  treatment  usually  consists  only  of  silicic  acid. 
To  prove  whether  it  consists  wholly  of  this  acid,  it  may  be  treated  with  H2SO4 
and  HF  as  described  in  P.  5. 

3.  Some  substances  which  are  not  much  acted  upon  by  alkali  carbonates 
alone  are  readily  attacked  when  an  oxidizing  substance  like  KNOs  is  present. 
Thus,  sulfides  are  converted  into  sulfates  and  chromium  compounds  (such  as 
chromite,  FeOC^Os)  into  chromates.     A  few  substances,  however,  such  as 
the  nativ  or  ignited  oxides  of  tin  and  aluminum,  may  be  only  partially  de- 
composed even  by  long-continued  fusion  with  the  mixed  fluxes.    Such  an  un- 
decomposed  residue  may  be  fused  with  KOH  in  a  nickel  or  silver  crucible  and 
the  fusion  treated  first  with  water  and  then  with  HC1.    The  oxides  of  aluminum 
and  stannic  tin,  if  finely  powdered,  dissolve  rapidly  in  fused  KOH.    The  aqueous 
extract  contains  the  aluminum  as  aluminate  and  most  of  the  tin  as  stannate. 
The  residue  undissolved  by  water  may  consist  of  black  nickel  oxide  from  the 
crucible,  stannic  hydroxide,  and  of  other  hydroxides  accompanying  the  aluminum 
or  tin  oxides;  all  of  which  dissolve  in  HC1. 

4.  A  few  milligrams  of  nickel  are  taken  up  from  the  crucible  by  the  flux,  so 
that  this  element,  as  well  as  the  alkali  elements,  cannot  be  tested  for  later  in 
the  analysis.     The  crucible  is,  however,  so  little  attacked  by  the  flux  that  it 
can  be  used  repeatedly. 

5.  Whenever  it  is  permissible,  it  is  somewhat  better  to  make  the  fusion  in 
a  platinum  crucible,  since  then  no  foreign  substances  are  introduced  from  the 
crucible.    It  is  not  permissible,  however,  to  ignite  in  platinum  vessels  compounds 
of  the  silver,  copper,  and  tin  groups;  for  these  may  be  reduced  to  the  metal 
by  heating  with  an  alkaline  flux.     The  same  is  true  of  sulfur,  sulfides,  and 
in  the  presence  of  organic  matter  of  phosphates;  for  all  these  elements  form 
easily  fusible  alloys  with  the  platinum,  and  thus  spoil  the  crucible.     More- 
over, alkaline  hydroxides  and  strongly  oxidizing  fluxes  (such  as  peroxides  and 
nitrates)  must  not  be  fused  in  platinum,  since  they  attack  it  fairly  rapidly. 
Therefore,  if  the  fusion  is  made  in  platinum,  no  more  KNOs  should  be  added 
than  is  necessary. 

Procedure  8. — Destruction  of  Organic  Matter. — If  the  closed-tube 
test  (P.  1)  has  shown  the  presence  of  organic  matter,  powder  or 
cut  into  small  pieces  1-5  g.  of  the  substance  (according  to  the  quan- 
tity of  organic  matter  present).  Add  to  it  in  a  casserole  about  5  cc. 
of  96%  H2S04;  warm  gently  until  the  substance  is  well  charred;  cool; 
add  slowly,  with  constant  stirring,  under  a  hood,  16-normal  HN03, 
until  violent  reaction  ceases;  warm  gently  for  a  few  minutes,  and 
then  heat  more  strongly,  keeping  the  dish  moving,  until  the  substance 
is  thoroly  charred.  Cool,  again  add  16-normal  HNOs  as  before,  and 

4 


40  PREPARATION  OF  THE  SOLUTION.  P.  8 

heat  until  thick  fumes  of  H2S04  are  evolved.  Repeat  this  process  till 
the  mixture  becomes  light-colored  and  remains  so  when  heated 
strongly. 

If  the  substance  has  dissolved  completely  (or  even  if  it  has  not,  if 
the  use  of  HF  is  impracticable),  evaporate  off  the  H2S04  under  a 
hood  till  only  1.5  cc.  remains,  cool  completely,  add  very  carefully 
10-20  cc.  water,  and  boil.  (If  a  precipitate  separates,  filter  it  off 
and  treat  it  by  P.  6.)  Treat  the  solution  by  P.  11. 

If  the  substance  has  not  dissolved  completely,  transfer  the  mixture 
to  a  platinum  crucible,  evaporate  off  the  £[2864  till  only  1.5  cc.  re- 
main, cool  completely,  and  treat  the  mixture  by  the  second  and 
third  paragraphs  of  P.  5. 

Notes. — 1.  This  method  of  destroying  organic  matter  is  of  very  general 
application,  being  effective  even  when  such  stable  substances  as  paraffin  and 
cellulose  are  present.  Organic  matter  can  also  be  destroyed  by  ignition,  but 
this  has  the  disadvantages  of  volatilizing  certain  elements,  especially  mercury 
and  arsenic,  and  of  making  some  substances  very  difficultly  soluble.  When 
the  organic  matter  consists  only  of  oil,  as  is  the  case  with  an  oil  paint,  it  may 
be  better  to  extract  it  with  ether,  especially  when  it  is  desired  to  determin  the 
proximate  constituents  of  the  substance. 

2.  The  residue  contains:   any  substances  originally  present  that  have  not 
been  attacked  by  HNOs  or  H2SO4,  especially  silicates;   all  the  lead,  strontium, 
and  barium  that  may  have  been  present  in  any  form,  since  the  sulfates  of 
these  elements  are  insoluble  in  dilute  H2SO4J   all  the  silica,  since  silicic  acid  is 
dehydrated  and  made  insoluble  by  heating  with  E^SC^;  some  of  the  calcium, 
bismuth,  antimony,  and  tin,  when  these  elements  are  present  in  considerable 
quantity,  since  their  sulfates  (or  oxides)  are  not  readily  soluble  in  dilute  H2SO4; 
and  substantially  all  of  the  chromium,  since  its  sulfate  is  converted  into  the 
insoluble  anhydrous  form. 

3.  After  the  organic  matter  is  destroyed,   the  solution  is  evaporated  to 
1.5  cc.,  in  order  that  the  concentration  of  the  acid  may  be  properly  adjusted 
in  the  subsequent  H2S  precipitation. 


DETECTION  OF  THE  BASIC  CONSTITUENTS. 


GENERAL   DISCUSSION. 

The  science  of  qualitativ  chemical  analysis  treats  of  the  methods 
of  determining  the  nature  of  the  elements  and  of  the  chemical  com- 
pounds which  are  present  in  any  given  substance.  When  the  presence 
or  absence  of  the  various  elements  is  alone  determined,  the  process  is 
called  ultimate  analysis;  when  the  chemical  compounds  of  which  the 
substance  is  composed  are  identified,  it  is  called  proximate  analysis.  In 
the  analysis  of  inorganic  substances,  to  which  this  book  is  devoted,  the 
object  in  view  is  ordinarily  an  intermediate  one — namely,  that  of  de- 
tecting the  base-forming  and  acid-forming  constituents  (called  in  this 
book  for  short  the  basic  and  acidic  constituents)  that  are  present  in  the 
substance.  Thus  the  analysis  of  a  substance  consisting  of  calcium  sul- 
fate,  zinc  chromate,  and  ferric  oxide  would  show  not  only  that  the  ele- 
ments calcium,  sulfur,  zinc,  chromium,  and  iron  were  present,  but  also 
that  the  sulfur  was  in  the  form  of  sulfate  (not  sulfide  or  sulfite),  the 
chromium  in  the  form  of  chromate  (not  of  a  chromic  salt),  and  the  iron 
in  the  ferric  (not  the  ferrous)  state.  The  reason  for  this  is  that  the 
analysis  is  carried  out  by  dissolving  the  substance  in  water  (with  aid 
of  acids,  if  necessary),  and  by  treating  the  solution  so  obtained  suc- 
cessively with  a  number  of  different  chemical  substances.  Now,  since 
the  chemical  reactions  of  substances  in  aqueous  solutions  are  deter- 
mined by  the  nature  of  the  ions  which  they  yield,  and  since  the  ions 
correspond  to  the  basic  and  acidic  constituents,  it  is  these  constitu- 
ents whose  presence  or  absence  is  established. 

For  detecting  the  basic  constituents  a  systematic  method  is  employed 
which  consists  in  adding  to  an  acid  solution  of  the  substance  in  suc- 
cession ammonium  chloride,  hydrogen  sulfide,  ammonium  hydroxide 
and  sulfide,  and  ammonium  carbonate.  By  each  of  these  reagents 
(which  is  the  name  applied  to  substances  added  to  produce  any  desired 
chemical  reaction)  a  group  of  basic  constituents  is  precipitated.  Thus, 
ammonium  chloride  precipitates  those  constituents  whose  chlorides 
are  only  slightly  soluble  in  water;  hydrogen  sulfide,  those  whose 
sulfides  are  only  slightly  soluble  in  dilute  acid;  ammonium  hydroxide 
and  sulfide,  those  whose  sulfides  or  hydroxides  are  only  slightly  soluble 
in  ammoniacal  solutions;  and  ammonium  carbonate,  those  whose  car- 
bonates are  only  slightly  soluble  in  water  containing  ammonium  car- 
bonate. The  way  in  which  the  basic  constituents  are  thus  separated 
into  groups  is  shown  in  more  detail  in  Table  II  on  the  following  page. 

41 


SEPARATION  INTO  GROUPS. 


HZS 


te. 


NHJ2S 


1 

i 


ni 


Filtrate: 
INUM-GRO 
CR  ZN) 


as  sodium  salts 
See  Table  VIII 


9  3  «o 

HI 

§§§ 

-I* 

^  2  '^ 

•*3    0!    3 

3  O  ' 

•'&  S  I 
SH^i 


...B|| 

a  §  oQ  S 

1^1 

^a^:! 


i 


ga- 

||| 

rt  g  ft 

O  cT 


..  fe'o 
l§^ 

•19^ 


p.  11 


ANALYSIS  OF  THE  SILVER-GROUP. 


43 


PRECIPITATION    AND   ANALYSIS    OF    THE    SILVER-GROUP. 


TABLE  III. — ANALYSIS  OF  THE  SILVER-GROUP. 


Precipitate:  BiOCl*,  PbCl2*,  AgCl,  Hg2Cl2.     Treat  with  HCl  (P.  12.) 


Solution:  BiCl3. 
Evaporate,  pour 
into  water 
(P.  IS). 

Residue:  PbCl2,  AgCl,  HgaCla.     Treat  with  hot  water  (P.  IS). 

Solution:  PbCl2. 
Add  HzSOt 
(P.  18). 

Residue:  AgCl,  Hg2Cl2. 
Pour  NH4OH  through  the  filter  (P.  14). 

Precipitate  : 
BiOCl. 

Black  residue: 
HgandHg^ 

Solution: 
Ag(NH3)2Cl. 
AddHNOz(P.14). 

Precipitate: 
PbS04. 

White  precipitate: 
AgCl. 

*  These  precipitates  form  only  when  large  quantities  of  bismuth  and  lead  are  present. 

Procedure  n. — Precipitation  of  the  Silver-Group. — Place  the  cold 
solution  of  the  substance  (prepared  by  P.  2,  3,  5,  6,  or  8,  and  containing 
5  cc.  6-normal  acid  in  about  25  cc.  of  solution)  in  a  conical  flask,  and 
add  to  it  10  cc.  NH4C1  solution.  (White  precipitate,  presence  of  SILVER- 
GROUP.)  Let  the  mixture  stand  for  3  or  4  minutes;  then  filter  it. 
(Precipitate,  P.  12;  filtrate,  P.  21.) 

Notes. — 1.  It  is  recommended  that  in  general  hard-glass  conical  flasks  (the 
to-called  Erlenmeyer  flasks  of  Jena  or  Bohemian  glass),  rather  than  beakers  or 
test-tubes,  be  employed  for  holding  solutions  that  are  being  subjected  to  the  opera- 
tions of  precipitation  and  heating. 

2.  Even  in  cases  where  it  is  not  essential  to  add  a  perfectly  deflnit  volume  of 
a  reagent,  the  analyst  should  make  it  a  practis  to  measure  out  the  quantity  to  be 
added,  rather  than  to  pour  in  an  indeflnit  quantity  from  the  reagent  bottle.  For 
this  purpose  a  10  cc.  graduate  should  be  constantly  at  hand.  For  adding  smaller 
quantities  than  1  cc.  a  dropper  should  be  used.  This  may  be  made  by  drawing 
out  one  end  of  a  short  glass  tube  to  a  wide  capillary  and  capping  the  other  end 
with  a  rubber  nipple.  When  more  of  a  reagent  than  is  needed  has  been  poured  into 
a  graduate  or  other  vessel,  it  should  never  be  poured  back  into  the  reagent  bottle, 
owing  to  the  danger  of  contaminating  the  reagent. 

5.  Unless  the  concentration  is  specified,  it  is  understood  that  all  salt  solutions 
used  as  reagents  are  1-normal,  that  is,  that  they  contain  one  equivalent  of  salt  per 
liter  of  solution',  also  that  the  acid  and  base  solutions  used  as  reagents  (those  of 
HCl,  HN03,  HzSOi,  HAc,  NH±OH,  and  NaOH)  are  6-normal. 

4.  By  one  equivalent  of  any  substance  is  meant  that  weight  of  it  which 
reacts  with  one  atomic  weight  (1.008  grams)  of  hydrogen  in  any  of  its  com- 
pounds or  with  the  weight  of  any  other  substance  which  itself  reacts  with  one 


44  ANALYSIS  OF  THE  SILVER-GROUP.  P.  11 

atomic  weight  of  hydrogen.  Thus,  one  equivalent  is  the  quantity  in  grams 
corresponding  to  the  following  formulas:  INaOH,  £Ba(OH)2,  1HC1,  £H2SO4, 
SH3PO4, 1NH4C1,  £Na2SO4,  |CaS04,  JFeCl3.  When  a  substance  may  take  part 
either  in  a  reaction  of  metathesis  or  in  one  of  oxidation  and  reduction,  its  meta- 
thetical  equivalent  has  to  be  distinguished  from  its  oxidation  equivalent.  Thus  the 
metathetical  equivalent  of  nitric  acid  is  1HNO3;  but  its  oxidation  equivalent 
(when  it  is  reduced  to  NO)  is  |HNO3.  In  this  book  the  term  equivalent  will 
always  be  used  to  denote  the  metathetical  equivalent. — Note  that  the  number 
of  equivalents  of  a  substance  is  a  certain  quantity  of  it;  but  that  the  term 
normal  denotes  its  concentration,  that  is,  the  quantity  of  it  per  unit-volume 
(more  specifically,  the  number  of  equivalents  of  it  per  liter). 

5.  If  NEUC1  produces  no  precipitate,  it  proves  the  absence  of  silver  and 
mercurous  mercury,  but  not  of  lead  or  bismuth,  since  PbCl2  is  fairly  soluble  in 
water  and  BiOCl  is  fairly  soluble  in  dilute  acid.     The  solubility  of  PbCl2   is, 
however,  much  less  in  a  solution  of  NH4C1  or  of  any  other  chloride  than  it  is  in 
water  owing  to  the  common-ion  effect  explained  in  detail  in  the  following  note. 

6.  The  fact  that  the  solubility  of  PbCl2  is  greatly  decreased  by  the  addition 
of  NH4C1  or  HC1  is  explained  as  follows:    The  mass-action  law  requires  that  at 
a  given  temperature  in  all  dilute  solutions  containing  lead  chloride  the  ratio  of 
the  product  of  the  ion-concentrations*  (Pb++)  X  (Cl~)2  to  the  concentration 
(PbCl2)  of  the  un-ionized  salt  have  the  same  value;  that  is,  (Pb++)  X  (Cl~)2  4- 
(PbCl2)  =  some  definit  value.     Now  in  all  solutions  which  have  been  saturated 
with  lead  chloride  as  a  result  of  sufficiently  long  contact  with  the  solid  substance, 
the  concentration  of  the  lead  chloride  present  as  such  (that  is,  as  un-ionized 
PbCl2)  must  evidently  have  the  same  value,  and  therefore  in  all  such  saturated 
solutions  the  ion-concentration  product  (Pb++)X(Cl~)2  must  also  have  the 
same  value;  that  is,  in  all  solutions  saturated  at  a  given  temperature  with  lead 
chloride,   (Pb++)  X  (Cl~)2  =  some  definit  value.     This  particular  value  which 
the  ion-concentration  product  has  when  the  solution  is  saturated  is  commonly 
called  the  solubility-product',  but.the  principles  involved  are  less  likely  to  be  mis- 
understood if  it  be  called  the  saturation-value  of  the  ion-concentration  product. 
The  saturation- value  varies,  of  course,  with  the  nature  of  the  salt,  and  with  the 
temperature  in  the  case  of  a  given  salt.     In  the  case  of  lead  chloride  at  20°,  whose 
solubility  in  water  at  20°  will  be  seen  by  reference  to  the  Table  on  page  122  to  be 
70  milli-equivalents  per  liter,  the  saturation-value  of  the  ion-concentration  prod- 
uct in  these  units  is  evidently  (70)  X  (70)  2  =  343, 000,  provided  the  ionization  be 
considered  to  be  complete,  as  may  be  assumed  to  be  true  in  these  qualitativ 
considerations  in  the  case  of  nearly  all  neutral  salts.     Any  solution  containing 
lead-ion  and  chloride-ion  in  which  the  ion-concentration  product  exceeds  this 
saturation-value  is  evidently  supersaturated  and  tends  to  deposit  the  solid  sub- 
stance;   and  any  solution  in  which  the  ion-concentration  product  is  less  than 
the  saturation-value  is  evidently  undersaturated  and  tends  to  dissolve  more 
of  the  solid  substance.    Now,  when  NH4C1  or  HC1  is  added  to  a  saturated  solu- 
tion of  PbCl2  in  water,  the  immediate  effect  is  to  increase  the  value  of  (Cl~), 
and  therefore  of  the  product    (Pb++)X(Cl~)2;    but    the    solution    becomes 
thereby  supersaturated,  and  PbCl2  will  precipitate  out  of  it  until  the  saturation- 
value  of  the  product  (Pb++)  X(CJ-)2  is  restored. 

*  In  mass-action  expressions  of  this  kind,  chemical  formulas  within  parentheses  denote  the 
concentrations  of  the  respectiv  substances,  that  is,  the  quantities  of  them  per  liter  of  solution. 


P.  12  ANALYSIS  OF  THE  SILVER-GROUP.  45 

Procedure  12. — Extraction  and  Detection  of  Bismuth. — Pour  re- 
peatedly through  the  filter  containing  the  NH4C1  precipitate  (P.  11) 
a  cold  10  cc.  portion  of  2-normal  HC1  (see  Note  1).  Treat  the 
residue  by  P.  13.  Evaporate  the  HC1  solution  in  a  casserole  almost  to 
dryness  (see  Note  2),  add  a  few  drops  of  water,  and  pour  the  solution 
into  a  flask  containing  100  cc.  water  previously  heated  to  50-70°. 
(Fine  white  precipitate,  presence  of  BISMUTH.)  Confirm  the  presence 
of  bismuth  by  filtering  off  the  precipitate  and  treating  it  by  P.  36. 

Notes. — 1.  When  it  is  directed  to  dissolve  a  precipitate  by  pouring  the  solvent 
repeatedly  through  the  filter,  this  is  best  done  by  pouring  a  single  portion  of  the  solvent 
from  one  test-tube  through  the  filter  into  another  test-tube  back  and  forth  three  or  four 
times.  When  the  solvent  is  to  be  used  hot  (as  in  P.  13),  it  should  be  heated  to  boiling 
between  each  pouring. 

2.  When  it  is  directed  to  evaporate  a  solution  almost  to  dryness  or  just  to  dryness, 
the  last  part  of  the  evaporation  should  be  carried  out  by  keeping  the  dish  moving  over 
a  small  flame  in  such  a  way  as  not  to  overheat  the  residue. 

3.  The  precipitate  is  extracted  with  2-normal  rather  than  with  6-normal 
HCl,  because  PbCl2,  AgCl,  and  Hg2Cl2  are  much  more  soluble  in  the  more  con- 
centrated acid,  owing  to  the  formation  of  acids  with  complex  anions,  such  as 
H+2PbCl4=,  H+2AgCl3-,  H+2HgCl4- 

4.  The  white  precipitate  of  BiOCl  formed  on  the  addition  of  water  in  the  test 
for  bismuth  is  produced  by  the  hydrolysis  of  Bids.     If  HCl,  the  other  product 
of  the  hydrolysis,  is  present  in  the  solution,  the  reaction  will  not  be  complete, 
and  a  greater  or  less  quantity  of  bismuth  will  remain  in  solution.     This  quantity 
increases  very  rapidly  with  the  acid  concentration.     For  this  reason,  nearly  all 
the  HCl  must  be  removed  by  evaporation  and  the  solution  must  be  added  to  a 
large  volume  of  water.     Warm  water  is  used,  because  the  precipitation  of  BiOCl 
takes  place  more  rapidly  at  the  higher  temperature.     Antimony  under  these 
conditions  givs  a  similar  precipitate;  but  it  can  not  be  present  in  the  HNOs 
solution  of  the  substance  in  quantity  sufficient  to  giv  a  precipitate  with  NH4C1, 
unless  the  substance  contained  also  much  chloride  or  sulfate.     If  it  were  so 
precipitated,  it  would  behave  in  P.  12  like  bismuth;  but  would  be  distinguished 
from  it  by  the  confirmatory  test  described  in  P.  36. 

5.  The  expression  for   the  solubility-product  of  BiOCl  is   (Bi+++)X(O)- 
X(C1~)=  const.     Acids  greatly  increase  its  solubility  because  their    H+    ion 
combines  with  the  O=  ion  to  form  water,  which  is  a  very  slightly  ionized  sub- 
stance.   Chlorides  also  increase  its  solubility;   for  the  increase  of  (Cl~)  which 
they  directly  produce  is  more  than  compensated  by  the  decrease  of  (Bi+++) 
which  arises  from  the  formation  of  un-ionized  Bids  (and  probably  also  of  com- 
plex anions,  such  as  BiCU"). 

Procedure  13. — Extraction  and  Detection  of  Lead. — Pour  repeatedly 
through  the  filter  containing  the  residue  undissolved  by  HCl  (P.  12) 
a  10  cc.  portion  of  boiling  water.  Wash  the  residue  thoroly  with 
hot  water,  and  treat  it  by  P.  14.  Cool  the  10  cc.  aqueous  extract 
and  add  to  it  10  cc.  H2SO4.  (Fine  white  precipitate,  presence  of  LEAD.) 
Confirm  the  presence  of  lead  by  filtering  out  the  precipitate,  washing 
it  with  a  little  water,  and  treating  it  by  P.  34. 


46  ANALYSIS  OF  THE  SILVER-GROUP.  P.  14 

Procedure  14. — Detection  of  Silver  and  Mercury. — Pour  repeat- 
edly through  the  filter  containing  the  residue  insoluble  in  hot  water 
(P.  13)  a  5-10  cc.  portion  of  NH4OH.  (Black  residue  on  the  filter, 
presence  of  MERCUROUS  MERCURY.)  Acidify  the  filtrate  with  HN03. 
(White  precipitate,  presence  of  SILVER.)  When  there  is  much  black 
residue  and  little  or  no  white  precipitate,  treat  the  residue  by  P.  15. 

Notes. — 1 .  When  two  quite  different  limiting  quantities  of  the  reagent  are  specified 
(for  example,  5—10  cc.  as  in  this  procedure),  the  quantity  added  should  be  adjusted 
to  the  size  of  the  precipitate. 

2.  The  black  residue  that  is  produced  by  the  action  of  NH4OH  on  H&Cl* 
is  a  mixture  of  finely  divided  mercury  with  the  white  mercuric  compound 
HgClNH2.    The  reaction  is  expressed  by  the  equation: 

Hg2Cl2+2NH4OH  =HgClNH2+Hg+NH4Cl+2H2O. 

The  compound  HgClNH2  may  be  considered  to  be  a  derivativ  of  HgCl2, 
formed  by  replacing  an  atom  of  chlorin  by  the  univalent  radical  NH2. 

3.  An  NH^OH  solution  contains  a  considerable  proportion  of  (unhydrated) 
NHs;  and  AgCl  dissolves  readily  in  it,  owing  to  the  formation  of  a  soluble  com- 
plex salt,  Ag(NH3)2Cl,  which  in  solution  is  largely  ionized  into    Ag(NH3)2"1" 
and  Cl~  ions.      This  complex  cation  has  so  slight  a  tendency  to  dissociate 
into   Ag+  and  NHs  that  the  ratio  of  its  concentration  to  that  of  the  simple 
Ag+  ion  is  about  107  in  a  normal  solution  of  NB^OH. 

4.  If  the  PbCl2  was  not  completely  extracted  from  the  NH4C1  precipitate 
by  boiling  water  (in  P.  13),  it  is  converted  into  a  basic  salt  (Pb(OH)Cl) 
by  the  NH4OH,  and  may  pass  through  the  filter,  yielding  a  turbid  filtrate. 
This  basic  salt  will,  however,  dissolve  on  the  addition  of  HNOa. 

Procedure  15. — Detection  of  Silver  in  the  Presence  of  Much  Mer- 
cury.— Wash  the  black  residue  undissolved  by  NH4OH  (P.  14),  and 
pour  repeatedly  through  the  filter  containing  it  a  mixture  of  3  cc.  HC1 
and  10  cc.  saturated  Br2  solution,  at  the  same  time  rubbing  the  residue 
with  a  glass  rod.  Wash  the  filter,  and  pour  repeatedly  through  it  a 
10  cc.  portion  of  NH4OH.  Acidify  the  solution  with  HN03.  (Yel- 
lowish-white precipitate,  presence  of  SILVER.) 

Notes. — 1.  When  much  mercury  is  present  a  considerable  quantity  of  silver 
(5  mg.  or  more)  may  be  so  completely  retained  in  the  black  residue  that  scarcely 
any  test  for  silver  is  obtained  in  P.  14.  This  is  probably  due  to  the  fact  that  the 
AgCl  is  reduced  to  metallic  silver  by  the  metallic  mercury.  When  much  mer- 
cury is  present  it  is  therefore  necessary  to  test  the  residue  for  silver,  as  described 
in  this  Procedure. 

2.  The  bromin  converts  the  mercury  in  the  residue  into  soluble  HgBr2  and 
the  silver  into  insoluble  AgBr.  The  HC1  dissolves  the  HgClNH2  present  in 
the  residue  with  formation  of  HgCl2. 


P.  SI  PRECIPITATION  OF  COPPER  AND  TIN  GROUPS.  47 


PRECIPITATION    AND    SEPARATION    OF   THE    COPPER    AND     TIN     GROUPS. 


TABLE  IV. — SEPARATION  OF  THE  COPPER  AND  TIN  GROUPS. 

HYDROGEN  SULPHIDE  PRECIPITATE:    HgS,  PbS,  61283,  CuS,  CdS; 

As2S3,  As2S5,  Sb2S3,  Sb2S6  SnS,  SnS2. 
Treat  with  ammonium  polysulfide  (P.  22). 


Residue:  HgS,  PbS, 
BijjSs,  CuS,  CdS. 
See  Table  V. 

Solution:  (NH4)3AsS4,  (NH4)3SbS4,  (NH4)2SnS3. 
Add  HCl  (P.  23). 

Precipitate: 
As2S5,  80285,  SnS2. 
See  Table  VI. 

Filtrate:  NH4C1. 
Reject. 

Procedure  21. — Precipitation  of  the  Copper  and  Tin  Groups. — 
Dilute  to  100  cc.  the  filtrate  from  the  NH4C1  precipitate  (P.  11) 
or  the  solution  of  the  substance  in  HCl  (P.  4),  which  should  contain 
just  5  cc.  of  6-normal  HN03,  H2S04,  or  HCl.  Place  this  solution 
in  a  conical  flask  provided  with  a  two-hole  rubber  stopper  in  which 
is  a  tube  leading  to  the  bottom  of  the  flask.  Pass  into  it  in  the  cold 
a  slow  current  of  H2S,  until,  upon  shutting  off  the  gas  and  shaking 
thoroly,  the  liquid  smells  strongly  of  H2S.  Filter  at  once,  wash 
the  precipitate  with  hot  water  (see  Note  1),  and  treat  it  by  P.  22. 
Heat  the  filtrate  nearly  to  boiling  (to. 70-90°),  and  pass  H2S  into  it 
at  that  temperature  for  5-10  minutes. 

If  there  is  no  further  precipitate,  treat  the  solution  by  P.  51. 

If  there  is  a  further  precipitate,  evaporate  the  mixture  almost  to 
dryness.  Add  3-5  cc.  12-normal  HCl  and  evaporate  just  to  dryness, 
to  destroy  the  HN03.  Then  add  10  cc.  6-normal  HCl,  saturate  the 
cold  solution  with  H2S,  heat  it  to  70-90°,  and  pass  H2S  into  it  for  5- 
10  minutes.  Cool  the  mixture,  dilute  it  to  100  cc.,  and  saturate  it 
with  H2S.  Filter  out  the  precipitate,  wash  it,  treat  it  by  P.  22,  and 
unite  the  residue  and  the  solution  with  those  coming  from  the  first 
H2S  precipitate.  Treat  the  filtrate  by  P.  51. 

Notes. — 1.  The  washing  of  precipitates  should  in  general  be  continued  until 
the  wash-water  will  no  longer  giv  a  test  for  any  substance  known  to  be  present  in 
the  filtrate  (for  example,  in  this  case  for  acid  with  blue  litmus-paper  or  for  chloride 
with  AgNO'S).  Precipitates  which  are  practically  insoluble  in  water  (like  all  the 
fulfides  and  hydroxides  that  are  met  with  in  this  System  of  Analysis)  are  best 
washed  with  nearly  boiling  water,  as  this  runs  through  the  filter  more  rapidly  and 
extracts  soluble  substances  more  readily.  Precipitates  which  are  appreciably 


48  PRECIPITATION  OF  COPPER  AND  TIN  GROUPS.  P.gl 

soluble  should  be  washed  with  cold  water  and  with  only  a  small  quantity  of  it.  The 
proper  method  of  washing  a  precipitate  is  to  cause  a  fine  stream  of  water  from  a 
wash-bottle  to  play  upon  the  upper  edge  of  the  filter.  The  wash-water  should  in 
general  not  be  allowed  to  run  into  the  filtrate,  so  as  not  to  dilute  it  unnecessarily. 
When,  however,  a  considerable  proportion  of  the  solution  is  likely  to  be  retained 
in  the  filter  and  precipitate,  it  is  well  to  add  the  first  washings  to  the  filtrate. 

2.  The  formation  of  a  white  precipitate  on  diluting  the  solution  to  100  cc. 
shows  the  presence  of  much  bismuth  or  antimony.    The  precipitate,  which  con- 
sists of  BiOCl  or  SbOCl,  need  not  be  filtered  off.    The  formation,  on  passing 
in  H2S,  of  a  white  or  yellowish  precipitate  which  immediately  turns  black  with 
more  H2S  indicates  mercury.     (The  white  compound  is  HgCl2.2HgS,  and  this 
is  converted  into  black  HgS  by  the  excess  of  H2S.)    The  formation  of  an  orange 
precipitate  shows  antimony;    of  a  yellow  one,  cadmium,  arsenic,  or  stannic 
tin.     All  the  other  sulfides  are  black. 

3.  The  solution  is  afterwards  heated  nearly  to  boiling  and  again  saturated 
with  H2S,  in  order  to  ensure  the  detection  of  arsenic;   for  this  element,  when 
present  in  the  higher  state  of  oxidation  (as  arsenic  acid)  is  only  very  slowly 
precipitated  by  H2S  in  the  cold.     At  70-90°  the  precipitation  is  much  more 
rapid,  especially  if  the  solution  has  been  previously  saturated  with  H2S  in  the 
cold.    Under  these  conditions  even  1  mg.  As  givs  a  distinct  precipitate  in  less 
than  5  minutes.    Continuous  treatment  with  H2S  at  70-90°  in  an  open  vessel 
does  not,  however,  completely  precipitate  a  large  quantity  of  arsenic  from 
such  a  weakly  acid  solution  even  within  an  hour.     For  this  reason,  when  a 
considerable  precipitate  forms  in  the  hot  solution,  it  is  directed  to  evaporate 
the  filtrate,  to  add  HC1  to  destroy  the  HNOs  (which  in  the  concentrated  state 
would  decompose  the  H2S),  to  dissolve  the  residue  in  HC1,  and  to  pass  H2S 
through  the  hot  solution.     From  this  concentrated  acid  solution  the  arsenic 
precipitates    completely   in    5-10    minutes.     The   reasons   for   this    peculiar 
behavior  of  arsenic  in  the  higher  state  of  oxidation  are  presented  in  Note  2, 
P.  43.     The  solution  is  finally  diluted  and  saturated  in  the  cold  with  H2S, 
since  the  other  elements  are  not  completely  precipitated  until  the  arsenic  has 
been  removed. 

4.  The  effect  of  acid  on  the  precipitation  of  the  sulfides  is  explained  by 
the  mass-action  law  and  ionic  theory  as  follows:  When  a  dilute  solution,  whether 
aqueous  or  acid,  is  saturated  at  a  definit   temperature  with  H2S  gas  under 
the  atmospheric  (or  any  definit)  pressure  the  H2S  present  as  such  always  hag 
the  same  concentration.    This  ionizes,  however,  to  a  slight  extent  into    H+ 
and   HS",   and  to   a  still  less  extent  into  2H+  and  S".    It  is  only  the  latter 
form  of  ionization  that  needs  to  be  considered  here.    Now  between  the  H2S 
and  its  ions  must  be  maintained  the  equilibrium  expressed  by  the   equation 
(H+)2X(S=)    =  const.  X  (H2S);  or,  since  in  this  case  (H2S)=  const.,  as  just 
stated,  it  follows  that  also  (H+)2X(S=)=  const.    From  this  it  is  evident  that 
when  (H+)  is  increased  by  the  addition  of  acid  to  the  solution,  (S*)   must  be 
decreased  in  the  proportion  in  which  the  square  of  (H+)  is   increased;   thus, 
if  (H+)   is  doubled,  (S=)  will  be  reduced  to  one-fourth.    But  in  order  that  a 
sulfide — for  example,  of  the  formula  M+^S" — may  precipitate,  the  concentra- 
tion-product (M++)  X  (S~)  must  attain  its  saturation-value.     This  value  varies, 
however,  with  the  nature  of  the  sulfide  and  with  the  temperature;   and  there- 
fore the  acid  concentration  that  will  barely  permit  of  precipitation  when  (M++) 


P.tl  PRECIPITATION  OF  COPPER  AND  TIN  GROUPS.  49 

has  a  definit  value  (for  example,  1  mg.  in  100  cc.)  will  be  different  for  different 
sulfides  and  for  the  same  sulfide  at  different  temperatures.  Thus  if  the  ele- 
ments are  arranged  in  the  order  in  which  they  are  precipitated  from  cold  HC1 
solutions  as  the  acid-concentration  is  progressivly  decreased,  the  series  is  ap- 
proximately as  follows:  arsenic,  mercury  and  copper,  antimony,  bismuth  and 
stannic  tin,  cadmium,  lead  and  stannous  tin,  zinc,  iron,  nickel  and  cobalt, 
manganese.  The  acid  concentration  which  permits  precipitation  also  varies 
with  the  ionization  of  the  acid;  thus  zinc  is  precipitated  from  a  fairly  concen- 
trated solution  of  acetic  acid,  since,  owing  to  the  slight  ionization  of  this  acid, 
the  H+  concentration  is  less  than  in  a  far  more  dilute  solution  of  HC1.  The 
three  acids,  HC1,  HNOs,  and  H2SC>4,  do  not,  however,  differ  greatly  in  this 
respect,  since  they  are  all  largely  ionized  in  dilute  solution. 

5.  The  acid  concentration  is  made  0.3  normal  (5  cc.  of  6-normal  acid  being 
present  in  100  cc.)  and  the  solution  is  saturated  with  H2S  gas  in  the  cold,  since 
under  these  conditions  even  1  mg.  of  cadmium,  lead,  or  tin  precipitates,  and 
even  300  mg.  of  zinc  remain  in  solution.     (This  statement  in  regard  to  zinc  is 
true,  however,  only  when  the  solution  contains  also  a  considerable  quantity 
of  chloride-ion,  such  as  was  added  in  P.  11,  and  when  it  is  not  allowed  to 
stand.)    Moreover,  even  when  a  small  quantity  of  any  of  the  iron-group  elements 
is  present  with  a  large  quantity  of  a  copper-group  element,  the  former  is  not 
carried  down  in  the  H2S  precipitate  under  these  conditions  to  such  an  extent 
as  to  prevent  its  detection  in  the  filtrate,  provided  as  much  of  it  as  1  mg.,  or  in 
some  combinations  as  much  as  2  mg..  is  present. 

6.  A  white,  finely  divided  precipitate  of  free  sulfur  will  be  formed  if  the 
solution  contains  substances  capable  of  oxidizing  H2S.    The  most  important 
of  these  likely  to  be  present  are  ferric  salts,  chromates,  permanganates,  and 
chlorates.    The  reduction  by  H2S  of  ferric  salts  to  ferrous  is  attended  by  a  change 
in  color  from  yellow  to  colorless;  of  chromates  to  chromic  salts,  from  orange  to 
green;    and  of  permanganates  to  manganous  salts,  from  purple  to  colorless. 
Nitric  acid,  if  it  were  fairly  concentrated,  would  also  destroy  the  H2S;  but  at 
the  concentration  in  question  (0.3  normal)  it  has  scarcely  any  oxidizing  action 
even  in  boiling  solution. 

7.  In  balancing  equations  expressing  reactions  of  oxidation  and  reduction, 
like  those  referred  to  in  the  preceding  note,  the  main  thing  is  to  determin  the 
number  of  molecules  of  the  oxidizing  and  reducing  substances  which  react  with 
one  another.    This  can  be  done  most  simply  by  considering,  in  the  way  illustrated 
by  the  following  examples,  the  changes  which  take  place  in  the  valences  of  the 
atoms  of  these  substances.     Thus,  in  the  reduction  of  a  ferric  to  a  ferrous  salt 
by  hydrogen  sulfide,  the  iron  atom  changes  in  valence  from  +3  to  +2,  and  the 
sulfur  atom  changes  in  valence  from  —2  (in  H2S)  to  zero  (in  ordinary  sulfur). 
Since  the  total  change  in  the  number  of  valences  must  be  equal  and  opposit 
in  the  two  substances,  it  is  evident  that  two  molecules  of  ferric  salt  react  with 
one  of  hydrogen  sulfide,  and  therefore  that  the  equation  is : 

2FeCl3 +H2S = FeCl2  +S  +2HC1. 

In  the  reduction  of  HClOs  to  HC1  by  H2S,  the  chlorin  atom  decreases  in  val- 
ence from  +5  to  —  1  (thus  by  six  positiv  valences) ;  hence  there  must  be  a  decrease 
of  six  negativ  valences  in  the  reducing  substance,  and  that  this  may  be  the  case 
three  molecules  of  H2S  are  evidently  required.  The  equation  is  therefore: 

HC103+3H2S =HC1+3S+3H20. 


50  SEPARATION  OF  COPPER  AND  TIN  GROUPS.  P.  22 

In  cases  where  the  valence  of  an  atom  in  a  substance  is  in  doubt,  it  can  be 
found  at  once  from  the  valences  of  the  other  atoms  with  the  aid  of  the  principle 
that  in  any  compound  the  sum  of  all  the  positiv  valences  is  equal  to  the  sum  of 
all  the  negativ  valences;  thus  in  chloric  acid  HClOs,  since  the  three  oxygen 
atoms  have  6  negativ  valences  and  the  hydrogen  atom  has  one  positiv  valence, 
the  chlorin  atom  must,  in  order  to  make  the  compound  neutral,  have  5  positiv 
valences. 

Consider  as  another  example  the  reduction  of  potassium  permanganate 
(KMnO*)  to  manganous  chloride  (MnCl2)  by  H2S  in  the  presence  of  HC1. 
The  number  of  valence  of  the  manganese  atom  is  seen  to  be  +7  in  KMn(>4 
(since  that  of  four  oxygen  atoms  is  —8  and  that  of  the  potassium  atom  is  +1) 
and  to  be  +2  in  MnCl2.  The  proportion  is  therefore  2KMnC>4:  5H2S,  and  the 
reaction  is: 

2KMnO4+5H2S+6HCl=2MnCl2+2KCl+5S+8H2O. 

The  amount  of  acid  required  in  such  cases  can  be  seen  by  inspection, — most 
readily  by  noting  how  many  hydrogen  atoms  are  needed  to  combine  with  the 
oxygen  atoms  of  the  substance  undergoing  reduction.  Thus  in  this  case  16 
hydrogen  atoms  are  evidently  needed  for  this  purpose;  and,  since  ten  are  fur- 
nished by  the  H2S,  six  more  must  be  supplied  by  adding  6  molecules  of  HC1 
(or  an  equivalent  quantity  of  some  other  acid). 

Procedure  22. — Separation  of  the  Copper  and  Tin  Groups  by  Am- 
monium Sulfide. — Transfer  the  H2S  precipitate  (P.  21)  to  a  small 
casserole,  add  to  it  10-25  cc.  ammonium  monosulfide  (if  the  original 
substance  was  treated  with  strong  HNOs  in  P.  3),  or  5-10  cc.  ammo- 
nium polysulfide  (if  it  was  dissolved  in  water  or  in  dilute  HNO* 
in  P.  2),  cover  the  dish,  and  warm  the  mixture  slightly  (to  40-60°) 
for  about  10  minutes  with  frequent  stirring.  Add  10  cc.  water,  filter, 
and  wash  once  with  hot  water.  [If  the  residue  is  large  and  much  has 
been  extracted  from  it  by  this  treatment,  as  indicated  by  its  appear- 
ance or  as  determined  in  P.  23,  warm  it  again  with  ammonium  mono- 
sulfide or  polysulfide,  and  filter,  collecting  the  filtrate  separate  from 
the  first  one.]  Wash  the  residue  thoroly  with  hot  water.  (Residue, 
P.  31;  solutions,  P.  23.) 

Notes. — 1.  When  a  small  precipitate  is  to  be  treated  with  a  solvent,  this  may  be 
done  by  pouring  a  portion  of  the  solvent  repeatedly  through  the  filter.  When  a 
considerable  precipitate  is  to  be  so  treated,  the  filter  is  opened,  the  portions  to  which 
no  precipitate  adheres  are  torn  off,  and  the  remainder  is  laid  along  the  side  of  a 
casserole;  the  solvent  is  then  poured  over  it  and  is  swashed  to  and  fro,  the  precipitate 
being  rubbed  at  the  same  time  with  a  glass  rod,  so  as  to  remove  it  from  the  filter.  If 
this  succeeds,  the  filter  is  drawn  out  and  thrown  away;  otherwise,  it  is  allowed  to 
disintegrate  and  filtered  out  together  with  any  residue. 

2.  The  ammonium  monosulfide  reagent  is  a  solution  of  (NBU^S  and  of  the 
products  of  its  hydrolysis,  NH^SH,  NH4OH,  and  a  little  H2S.  The  polysulfide 
reagent  contains  in  addition  some  of  the  disulfide  (NH4)2S2  and  of  its  hydrolysis 
product  (NH4)HS2.  The  monosulfide  is  prepared  by  saturating  concentrated 


P.8S  SEPARATION  OF  COPPER  AND  TIN  GROUPS.  51 

NH4OH  with  H2S  gas  (whereby  NH4SH  is  formed),  adding  an  equal  volume 
of  NH4OH,  and  diluting  the  solution.  The  polysulfide  is  prepared  from  this 
solution  by  dissolving  sulfur  in  it.  These  reagents,  especially  the  monosul- 
fide,  should  be  kept  as  far  as  possible  out  of  contact  with  the  air,  which  is  con- 
veniently done  by  storing  them  in  small,  completely  filled,  glass-stoppered 
bottles;  for  the  oxygen  of  the  ah*  destroys  the  sulfide  with  liberation  of  sulfur, 
which  at  first  combines  with  the  still  unchanged  sulfide  forming  the  polysulfide, 
but  later  precipitates  when  the  oxidation  becomes  more  complete. 

3.  The  chemical  action  of  ammonium  sulfide  in  dissolving  the  sulfides  of 
the  tin-group  depends  on  the  formation  of  soluble  salts  of  sulfoacids  with 
complex  anions.  When  the  monosulfide  is  used,  the  reactions  are  as  follows  : 

_     j  (NH4+)3AsS3 


(  AsoSs  )  +3(NH4)2S  _ 
H 


+  (NH4)2S  =       (NH4+)2SnSr. 

The  disulfide  present  in  the  polysulfide  oxidizes  the  lower  sulfides  (As2S3,  SbjjSt, 
SnS)  to  the  same  sulfosalts  as  are  obtained  by  dissolving  the  higher  sulfides 
(A82S5,  80285,  SnS2)  in  ammonium  monosulfide.  It  will  be  seen  that  these  sul- 
fo-salts  are  analogous  to  the  salts  of  the  familiar  oxygen  acids  of  these  ele- 
ments, the  difference  being  that  sulfur  has  replaced  oxygen;  and  they  are  so 
named  as  to  indicate  this  relationship.  Thus  the  five  sulfo-salts  whose  formulas 
are  given  above  are  called  ammonium  sulfarsenite,  sulfantimonite,  sulfarsenate, 
sulfantimonate,  and  sulfostannate. 

4.  In  separating  the  copper-group  from  the  tin-group  (colorless)  ammonium 
monosulfide  is  used  rather  than  (yellow)  polysulfide  whenever  the  H2S  pre- 
cipitate must  contain  in  the  state  of  the  higher  sulfide  (81182  or  Sb2S5)  any  tin 
or  any  antimony  that  is  present.     This  is  the  cas3  when  hot  concentrated  nitric 
acid  was  used  originally  in  dissolving  the  substance,  but  may  not  be  so  when 
water  or  dilute  HNO3  was  used;  hence  the  directions  as  to  the  choice  between 
the  two  solvents.     The  polysulfide  has  the  disadvantages  that  it  yields  a  large 
precipitate  of  sulfur  on  acidification  and  that  it  dissolves  a  not  inconsiderable 
quantity  of  CuS  and  HgS,  thus  making  it  more  difficult  to  determin  from  the 
color  of  the  HC1  precipitate  obtained  in  P.  23  whether  or  not  elements  of  the 
tin-group  are  present.     The  fact  that  it  dissolves  some  CuS  and  HgS  also 
diminishes  the  delicacy  of  the  tests  for  copper  and  mercury.     The  polysulfide 
must,  nevertheless,  be  used  if  tin  may  be  present  as  SnS,  or  much  antimony 
as  Sb2Ss;  for  in  the  monosulfide  SnS  is  only  slightly  soluble,  and  Sb2S3  only 
moderately  soluble. 

5.  More  specifically,  the  behavior  of  the  various  sulfides,  when  warmed  with 
10  cc.  of  these  reagents,  is  as  follows:  Of  the  sulfides  of  the  copper-group  none 
dissolves  to  a  significant  extent  in  ammonium  monosulfide.     5-10  mg.  CuS  and 
0.5-1.0  mg.  HgS  may,  however,  dissolve  in  the  polysulfide  when  the  substance 
contains  a  large  quantity  of  these  elements.    Yet  when  only  2  mg.  are  present, 
either  of  these  elements  can  be  detected  in  the  analysis  of  the  copper-group, 
even  when  the  polysulfide  is  used,  provided  only  one  treatment  with  it  has 
been  made.    Of  the  sulfides  of  the  tin-group,  500  mg.  of  As  as  A^Sa  or  As2Ss, 
of  Sb  as  Sb2S5,  or  of  Sn  as  SnS2  dissolves  in  either  the  monosulfide  or  polysul- 


52  SEPARATION  OF  COPPER  AND  TIN  GROUPS.  P.  £S 


fide,  and  500  mg.  of  Sn  as  SnS  or  of  Sb  as  Sb2Sa  dissolve  in  the  polysulfide. 
Scarcely  any  SnS  and  only  50-100  mg.  Sb  as  80283  dissolve  in  the 
monosulfide. 

6.  Even  when  a  quantity  of  only  1  or  2  mg.  of  arsenic  or  antimony  is  present 
with  a  large  quantity  (even  500  mg.)  of  an  element  of  the  copper-group,  enough 
is  extracted  by  either  the  monosulfide  or  polysulfide  to  be  detected  in  the  sub- 
sequent tests.  With  tin,  however,  the  separation  is  imperfect;  for,  when  a 
large  quantity  of  elements  of  the  copper-group  and  only  3-5  mg.  of  tin  are 
present,  the  whole  of  this  may  remain  undissolved;  indeed,  when  much  cad- 
mium is  present  and  the  tin  is  in  the  stannous  state,  as  much  as  15  mg.  of  the 
latter  may  be  wholly  left  in  the  residue,  even  when  the  polysulfide  is  used. 
On  this  account  it  is  necessary  to  test  for  tin  in  the  course  of  the  analysis  of 
the  copper-group. 

Procedure  23.  —  Reprecipitation  of  the  Tin-Group.  —  Dilute  in  a 
small  flask  the  first  portion  of  the  ammonium  sulfide  solution  (P.  22) 
with  about  20  cc.  water,  make  it  slightly  acid  with  HC1  (see  Note  2), 
and  shake  the  mixture  for  2  or  3  minutes  to  coagulate  the  precipitate. 
(White  or  pale  yellow  precipitate,  absence  of  TIN-GROUP;  deep  yellow 
or  orange  precipitate,  presence  of  TIN-GROUP.  —  See  Notes  4  and  5.) 
[Treat  the  second  portion  of  the  ammonium  sulfide  solution  (P.  22) 
in  the  same  way,  and  unite  the  precipitate,  if  considerable  in  amount, 
with  the  first  one.]  Filter  out  and  wash  the  precipitate,  using  suction, 
finally  sucking  it  as  dry  as  possible.  (Precipitate,  P.  41;  filtrates, 
reject.) 

Notes.  —  1.  In  cases  where  the  filtration  is  slow,  where  the  precipitate  must  be 
washed  with  very  little  water,  or  where  (as  in  this  case}  it  must  be  free  '  as  far 
as  possible  from  water,  it  is  advisable  to  filter  with  the  aid  of  suction.  This  op- 
eration is  carried  out  by  reinforcing  the  ordinary  filter  with  a  small  hardened 
filter  placed  below  it  in  the  funnel,  by  inserting  the  funnel  in  a  rubber  stopper  in 
the  neck  of  a  filter-bottle,  and  connecting  the  side  arm  of  the  filter-bottle  to  a 
suction-pump  by  means  of  a  rubber  tube  carrying  a  screw-clamp.  The  suction 
should  be  applied  very  gradually  so  as  to  avoid  breaking  the  filter.  The  filtrate 
should  be  poured  out  of  the  filter-bottle  before  beginning  to  wash  the  precipitate. 

2.  Whenever  it  is  directed  to  make  a  solution  slightly  acid  or  alkaline,  this 
should  be  done  carefully  as  follows:    Add  from  a  graduate  somewhat  less  acid  or 
alkali  than  will  neutralize  the  alkali  or  acid  known  to  be  present  in  the  solution. 
Then  add  from  a  dropper  more  of  the  acid  or  alkali,  a  drop  or  two  at  a  time,  till 
a  glass  rod  dipped  in  the  solution  and  touched  to  a  piece  of  blue  or  red  litmus  paper 
placed  on  a  watch-glass  changes  the  color  of  the  litmus. 

3.  When  the  HC1  is  added  to  the  solution  of  the  sulfosalts,  the  correspond- 
ing sulfoacids  which  are  liberated  decompose  immediately  into  H2S  and  the 
solid  sulfides.    These  are  now  necessarily  in  the  higher  state  of  oxidation,  since 
the  lower  sulfides,  if  originally  present,  have  been  oxidized  by  the  polysulfide. 
The  fact  that  the  sulfide  precipitates  when  the  solution  of  the  sulfosalt  is  acidified 
is  a  consequence  of  the  mass-action-law.     Thus,  since  the  complex  anions 
dissociate  according  to  the  equations, 

80285+  3S~, 


P.  SS  SEPARATION  OF  COPPER  AND   TIN  GROUPS.  53 

this  law  evidently  requires  that,  in  any  solution  saturated  with  the  solid  sulfide, 
the  concentration  of  the  complex  anion,  and  therefore  of  the  tin,  arsenic,  or 
antimony  in  the  solution,  increase  with  increasing  concentration  of  the  S"  ion. 
Now  in  the  solution  of  the  largely  ionized  (NEU^S  there  is  a  large  concentration 
of  S"  ion;  but  when  the  solution  is  made  acid  with  HC1,  the  S"  ion  is  almost 
completely  converted  by  the  relativly  large  concentration  of  the  H+  ion  into  the 
slightly  ionized  substances  HS~  and  H2S. 

4.  Much  time  is  ordinarily  saved  by  determining  at  this  point  whether  or 
not  any  element  of  the  tin-group  is  present.    When  ammonium  monosulfide 
has  been  used,  there  is  never  any  difficulty  in  drawing  a  definit  conclusion 
in  regard  to  this  from  the  size  and  appearance  of  the  HC1  precipitate;   for 
in  the  absence  of  the  tin-group  only  a  very  small,  nearly  white  precipitate  of 
finely  divided  sulfur  separates.     Even  when  ammonium  poly  sulfide  has  been 
used,    it    is    usually    possible  to  decide  as   to   the  presence   or   absence   of 
a  small  quantity  of  arsenic,  antimony,  or  tin;   for,  altho  considerable  sulfur  is 
then  liberated,  yet  the  precipitate  is  distinctly  lighter  colored  and  smaller  in 
appearance  when  it  consists  only  of  sulfur  than  it  is  when  even  1  mg.  of  arsenic, 
antimony,  or  tin  is  present  with  the  sulfur.    The  difference  is  most  marked  when 
only  a  slight  excess  of  HC1  has  been  used  in  the  acidification  and  when  the 
acidified  mixture  has  been  shaken,  but  not  heated.     In  any  doubtful  case,  the 
precipitate  should  be  compared  with  that  produced  by  acidifying  5-10  cc.  of  am- 
monium polysulfide  to  which  1  mg.  of  arsenic,  antimony,  or  tin  has  been  added. 

5.  When,  however,  the  HC1  precipitate  from  a  polysulfide  solution  is  fairly 
small  and  is  dark  brown  (indicating  copper)  or  dark  gray  or  black  (indicating 
mercury)  or  of  unpronounced  yellow  or  orange  color,  so  as  to  make  any  con- 
clusion as  to  the  tin-group  doubtful,  the  precipitate  is  best  treated  as  follows: 
Heat  it  with  15-20  cc.  NEUOH  almost  to  boiling  for  5  minutes  and  filter; 
test  the  precipitate  for  copper  by  P.  31,  35,  and  37,  if  it  has  not  already  been 
found  present;    add  to  the  filtrate  a  few  drops  of  ammonium  monosulfide, 
filter  out  any  precipitate,  heat  the  filtrate  to  boiling,  make  it  acid  with  HC1, 
shake,  filter  out  the  precipitate,  and  treat  it  by  P.  41  as  usual.    The  character 
of  the  HC1  precipitate  now  obtained  will  clearly  indicate  the  presence  or  ab- 
sence of  the  tin-group;  for  by  the  treatment  with  NELjOH  the  excess  of  sulfur 
originally  present  and  any  CuS  is  left  undissolved,  and  by  the  (NH^S  added 
to  the  solution  any  mercury  present  is  precipitated,  so  that  the  HC1  precipi- 
tate can  contain  only  sulfides  of  the  tin-group  and  a  very  little  sulfur.    A^Ss, 
Sb2S5,  and  SnS2  all  dissolve  in  NEUOH  (tho  in  the  cases  of  Sl^Ss  and  SnSj 
less  abundantly  than  in  ammonium  sulfide),  owing  to  the  formation  of  a  mix- 
ture of  salts  of  partially  sulfurated  acids,  such  as  HsAsOgS  and  HsAsC^. 
The  addition  of  (NIL^S  to  the  NEUOH  solution  and  the  heating  serve  to 
convert  these  into  the  fully  sulfurated  acids,  such  as  HgAsS4  ;   from  which  HC1 
will  then  precipitate  the  simple  sulfides  much  more  completely. — The  inci- 
dental removal  of  the  small  amounts  of  CuS  and  HgS  by  the  NKUOH  treat- 
ment is  not  necessary  so  far  as  the  analysis  of  the  tin-group  is  concerned,  since 
their  presence  does  not  interfere  with  the  detection  of  even  1  mg.  of  arsenic, 
antimony,  or  tin;  but  it  does  enable  1  or  2  mg.  of  copper  to  be  detected  which 
might  otherwise  be  lost. — In  this  connection  it  may  be  mentioned  that  a  mix- 
ture of  SnS2  and  SbzS$  does  not  always  have  a  color  intermediate  between  the 
yellow  and  orange  colors  of  the  separate  sulfides,  but  that  it  may  be  brown  or 
dark  gray. 


54 


ANALYSIS  OF  THE  COPPER-GROUP. 
ANALYSIS  OF  THE  COPPER-GROUP. 


P.  31 


TABLE  V. — ANALYSIS  OF  THE  COPPER-GROUP. 


RESIDUE  FROM  AMMONIUM  SULFIDE  TREATMENT!    HgS,  PbS,  61283,  CuS,  CdS. 

Boil  vrilh  HN03  (P.  81). 


Residue:  HgS. 
Add  Br2  solution  (P.  82}  . 

Solution:  Pb,  Bi,  Cu,  Cd  as  nitrates. 
Add  HySOt,  evaporate,  add  water  (P.  55). 

Residue: 
Sulfur. 

Solution:  HgBr2. 
Add  SnCk- 

Precipitate: 
PbS04. 
Dissolve  in 
NHiAc, 
add  K2Cr04 
(P.  84). 

Filtrate.    Add  NH^OH  (P.  85}  . 

Precipitate: 
Bi(OH)3. 
Add  NazSn02. 
(P.  36). 

Filtrate:  Cu(NH3)4SO4, 
Cd(NH3)4S04. 

White  or  gray 
precipitate: 
Hg2Cl2  or  Hg. 

To  a  small 
part  add 
H  Ac  and 

KtFe(CN)t 
(P.  57). 

To  the 
remainder 
addKCN 
and  H2S 
(P.  88). 

Yellow 
precipitate: 
PbCr04. 

Black  residue: 
Bi. 

Red 
precipitate: 
Cu2Fe(CN)* 
White 
precipitate  : 
Cd2Fe(CN)6. 

Yellow 
precipitate: 
CdS. 
Solution: 
K3Cu(CN)4. 

Procedure  31. — Treatment  of  the  Sulfides  with  Nitric  Add. — To 
the  residue  from  the  ammonium  sulfide  treatment  (P.  22)  in  a  casse- 
role add  10-20  cc.  2-normal  HNO3  solution,  heat  to  boiling,  and  boil 
gently  for  a  minute  or  two.  Filter,  and  wash  the  residue.  (Residue, 
P.  32;  solution,  P.  33.) 

Notes. — 1.  Boiling  HNOs  of  this  concentration  dissolves  the  sulfides  of 
lead,  bismuth,  copper,  and  cadmium  almost  immediately,  and  is  therefore 
preferable  to  a  more  dilute  acid,  with  which  the  reaction  would  require  for 
its  completion  several  minutes'  boiling.  Only  a  little  HgS  is  dissolved  by 
this  treatment,  unless  the  boiling  is  long  continued. 

2.  Moderately  concentrated  H&Os  dissolves  sulfides  much  more  rapidly 
than  HC1  or  H2SO4  of  the  same  concentration;  for  with  the  latter  acids  the 
sulfide-ion  is  removed  from  the  solution  only  by  combination  with  the  hydrogen- 
ion  and  by  the  volatilization  of  the  H2S  formed  thereby,  while  with  HNOs 
the  sulfide-ion  (or  the  H2S  in  equilibrium  with  it)  may  also  be  destroyed  by 
oxidation  to  ordinary  sulfur.  The  oxidizing  action  of  HNOa  is,  however,  slow, 
unless  it  is  hot  and  as  concentrated  as  2-normal. 


P.  SI  ANALYSIS  OF  THE  COPPER-GROUP.  55 

3.  That  HgS,  unlike  the  other  sulfides,  does  not  dissolve  in  the  dilute  HNOs 
is  doubtless  due  to  the  much  smaller  concentration  of  its  ions  in  its  saturated 
solution  and  to  the  fact  that  at  this  small  concentration  sulfide-ion  (or  the 
£[28  in  equilibrium  with  it  at  a  correspondingly  small  concentration)  is  oxidized 
only  very  slowly  by  the  dilute  HNOa.    HgS  is,  however,  readily  dissolved  by 
more  vigorous  oxidizing  agents,  such  as  aqua  regia  or  bromin  solution,  since  they 
react  rapidly  with  sulfide-ion  (or  with  H2S)  even  when  its  concentration  is  very 
small. 

4.  If  more  concentrated  HNOa  be  used,  or  if  the  acid  become  concentrated 
by  long  boiling,  the  black  HgS  is  dissolved  in  part,  and  the  remainder  is  con- 
verted into  a  heavy,  white,  difficultly  soluble  compound  (Hg(NC>3)2.2HgS).    As 
a  small  quantity  of  HgS  may  be  so  converted  even  by  the  hot  2-normal  HNOa, 
the  residue  should  be  tested  for  mercury  by  P.  32,  even  if  it  is  light-colored. 
A  light-colored  residue  must  often  be  tested  also  for  tin  (see  Notes  2  and  3, 
P.  32). 

5.  When  much  lead,   copper,   or  bismuth  is  present  the  sulfur  formed 
generally  encloses  enough  of  the  undissolved  sulfide  to  give  it  a  black  color. 
A  black  residue  is  therefore  not  necessarily  HgS,  but  must  be  further  tested  for 
mercury  as  described  in  P.  32. 

6.  Some  sulfur  is  always  oxidized  to  H2SO4  by  the  boiling  HNOs;    but, 
even  in  the  presence  of  much  lead,  PbSO4  is  not  precipitated,  owing  to  its 
moderate  solubility  in  HNOs. 

7.  Even  when  only  1  mg.  of  lead,  bismuth,  or  copper  is  present  with  500  mg. 
of  mercury,  enough  of  any  of  these  elements  is  extracted  from  the  residue  by 
the  HNOa  to  giv  a  satisfactory  test  in  the  subsequent  procedures;  but  in  the 
case  that  1—5  mg.  of  cadmium  are  associated  with  300-500  mg.  of  mercury, 
about  four-fifths  of  the  cadmium  is  retained  in  the  residue,  so  that  a  corre- 
spondingly small  precipitate  of  CdS  is  produced  in  the  final  test  for  cadmium, 
and  1-2  mg.  may  escape  detection. 

Procedure  32. — Confirmatory  Test  for  Mercury. — Transfer  the 
residue  undissolved  by  HNOs  (P.  31),  with  the  filter  if  necessary,  to 
a  casserole,  add  10-40  cc.  saturated  Br2  solution,  cover  the  dish,  and 
warm  slightly  for  5-10  minutes,  with  frequent  stirring.  Boil  the 
mixture  until  the  bromin  is  expelled,  and  filter.  (Residue,  see  Note  3.) 
Cool  the  solution,  and  add  to  it  15-20  drops  of  HC1  and  one  drop  of 
SnCU  solution;  then  add  3-5  cc.  SnCl2  solution.  (White  pre- 
cipitate turning  gray,  or  gray  precipitate,  presence  of  MERCURY.) 

Notes. — 1.  In  the  final  test  for  mercury  HC1  is  added  to  prevent  the  pre- 
cipitation  of  a  basic  tin  salt  when  the  SnCl2  reagent  is  diluted,  to  diminish 
the  tendency  of  CuCl  to  precipitate, "and  to  cause  the  formation  at  first  of 
white  Hg2Cl2.  For  the  last  reason,  also,  a  single  drop  of  SnCk  solution  is  first 
added  to  the  cold  solution.  By  the  excess  of  SnCfe  the  white  precipitate  is  re- 
duced to  gray,  finely  divided  mercury.  This  darkening  of  the  precipitate  dis- 
tinguishes Hg2Cl2  from  CuCl,  which  may  separate  as  a  white  precipitate  if  a 
large  quantity  of  CuS  was  left  undissolved  by  the  HNOs  in  P.  31. 

2.    If   elements  of   the  copper-group  are  present  in  large  quantity  (50- 


56  ANALYSIS  OF  THE  COPPER-GROUP.  P.  32 

500  mg.),  the  residue  from  the  Br2  treatment  should  be  tested  for  tin,  in  order 
to  guard  against  overlooking  the  presence  of  a  small  quantity  of  this  element 
in  the  substance.  For,  as  stated  in  P.  22,  Note  6,  a  quantity  of  tin  as  large 
as  5  mg.  (or  even  larger  when  stannous  tin  and  cadmium  are  both  present) 
may  remain  entirely  in  the  residue  undissolved  by  ammonium  sulfide  when 
this  residue  is  large.  Any  SnS  or  SnS2  present  in  it  is  converted  by  the  HNOs 
in  P.  31  into  metastannic  acid  (H2SnOs),  very  little  of  which  is  dissolved 
either  by  the  acid  or  by  the  Br2  solution  used  in  P.  32. 

3.  To  recover  the  tin  from  the  residue,  proceed  as  follows :  Digest  the  residue 
from  the  Br2  treatment,  if  it  is  still  dark-colored,  with  another  portion  of  Br2 
solution  (to  extract  the  rest  of  the  mercury),  filter,  reject  the  filtrate,  and  warm 
the  residue  slightly  with  2-5  cc.  ammonium  monosulfide.    Filter,  and  unite  the 
solution  with  the  main  ammonium  sulfide  solution  obtained  in  P.  22.    By  this 
procedure  2  mg.  Sn  can  be  detected. 

4.  In  case  tin  need  not  be  tested  for  at  this  point,  the  residue  may  be  more 
quickly  dissolved  by  warming  it  with  HC1  and  adding  gradually  a  little  solid 
KClOs.    Metastannic  acid,  being  soluble  in  HC1,  then  passes  into  solution  with 
the  mercuric  salt,  but  it  does  not  interfere  with  the  test  for  mercury. 

Procedure  33. — Precipitation  of  Lead  with  Sulfuric  Acid. — To  the 
HN03  solution  (P.  31)  add  10  cc.  H2S04  and  evaporate  in  a  casserole 
to  a  volume  of  1-2  cc.  until  dense  white  fumes  of  H2S04  begin  to 
come  off.  Cool  the  mixture  and  pour  it  into  10-15  cc.  cold  water, 
rinsing  out  the  casserole  with  the  same  solution.  Cool  again,  shake, 
and  allow  the  mixture  to  stand  5  minutes,  but  not  much  longer. 
(Fine  white  precipitate,  presence  of  LEAD.)  Filter,  and  wash  the  pre- 
cipitate first  with  H2S04,  and  finally  with  a  little  water.  (Precipitate, 
P.  34;  filtrate,  P.  35.) 

Notes. — 1.  PbSC>4  is  somewhat  soluble  both  in  water  and  in  concentrated 
H2SO4,  but  much  less  so  in  dilute  H2SO4,  its  solubility  being  scarcely  appreciable 
in  the  cold  6-normal  acid.  That  the  solubility  in  dilute  H2SO4  is  less  than  that 
in  water  is  due  mainly  to  the  common-ion  effect.  Concentrated  H2SC>4  is,  of 
course,  an  entirely  different  solvent.  PbSO4  dissolves  fairly  readily  in  dilute 
HNOs,  owing  to  the  tendency  to  form  the  intermediate  HSO4~  ion;  hence,  to 
ensure  complete  precipitation  of  PbS04,  the  HNOs  must  be  removed  by 
evaporation.  To  ensure  complete  precipitation  it  is  also  necessary  to  let  the 
mixture  stand  a  few  minutes;  for  the  solutions  of  substances  which,  like  PbSO4, 
are  crystalline  and  appreciably  soluble  tend  to  remain  supersaturated. 

2.  When  much  bismuth  is  present  it  ordinarily  dissolves  at  first  when  the 
water  is  added  to  the  concentrated  H2SO4,  provided  the  mixture  is  kept  cold; 
but  from  this  solution  a  coarsely  crystalline  precipitate  of  an  oxysulfate,  such 
as  (BiO)2S04,  separates  slowly  upon  standing  in  the  cold  but  almost  imme- 
1  diately  upon  heating,  and  to  such  an  extent  that  there  may  remain  in  solution 
not  more  than  50  mg.  of  bismuth.  If  the  precipitate  is  of  this  character,  free 
it  from  bismuth  before  applying  the  confirmatory  test  for  lead  by  pouring 
repeatedly  through  the  filter  a  5-10  cc.  portion  of  HC1  and  treating  the  solution 
so  obtained  by  P.  33.  The  evaporation  with  H2S04  is  necessary  in  order  to 
ensure  reprecipitation  of  the  PbSO4  that  has  been  dissolved  by  the  HC1. 


P.  34  ANALYSIS  OF  THE  COPPER-GROUP.  57 

Procedure  34. — Confirmatory  Test  for  Lead. — Pour  repeatedly 
through  the  filter  containing  the  H2S04  precipitate  (P.  33)  a  10-20 
cc.  portion  of  NH4Ac  solution.  (See  Note  1,  P.  12.)  To  the  filtrate 
add  a  few  drops  of  K2CrO4  solution  and  2-5  cc.  HAc.  (Yellow  precipi- 
tate, presence  of  LEAD.) 

Notes. — 1.  This  confirmatory  test  for  lead  should  not  be  omitted;  for  the 
HjS04  precipitate  may  consist  not  only  of  PbSO*  but  of  (BiO)2SO4  or  of  BaSC>4, 
which  last  closely  resembles  PbS(>4  in  appearance.  (BiO)2SO4  dissolves  in  NH^Ac 
solution  and  gives  a  yellow  precipitate  on  adding  K2OO4 ;  but  this  precipitate, 
unlike  PbCrO*,  dissolves  readily  in  acetic  acid. 

2.  The  solubility  of  PbSO4  in  NH4Ac  solution  depends  on  the  formation 
by  metathesis  of  undissociated  PbAcs,  this  salt  being  much  less  ionized  than 
most  other  salts  of  the  same  valence-type.  On  the  addition  of  a  chromate  to 
this  solution  the  much  more  difficultly  soluble  PbO04  is  precipitated.  BaSC>4 
is  not  dissolved  by  NEUAc  solution,  owing  to  its  very  slight  solubility  in  water 
and  the  fact  that  barium  acetate,  unlike  lead  acetate,  is  a  largely  ionized  salt. 

Procedure  35. — Precipitation  of  Bismuth  with  Ammonium  Hy- 
droxide—-To  the  H2S04  solution  (P.  33)  add  NH4OH  slowly  until 
a  strong  odor  of  it  persists  after  shaking.  (White  precipitate,  possible 
presence  of  BISMUTH;  blue  solution,  presence  of  COPPER.)  Shake  to 
cause  coagulation,  filter,  and  wash  the  precipitate  thoroly.  (Precipi- 
tate, P.  36;  filtrate,  P.  37  and  38.) 

Notes. — 1.  The  precipitate  produced  by  NH4OH  may  consist  also  of  Fe(OH)a, 
or  of  other  hydroxides  of  the  iron-group,  if  these  elements  were  carried  down 
in  the  H^S  precipitate  or  were  not  completely  removed  from  it  by  washing. 
The  formation  of  a  small  precipitate  is,  therefore,  not  a  sufficient  proof  of  the 
presence  of  bismuth,  and  the  confirmatory  test  of  P.  36  must  be  applied. 

2.  Cd(OH)2  or  Cu(OH)2,  tho  only  very  slightly  soluble  in  water,  dissolves 
in  NH^OH  owing  to  the  combination  of  the  Cd++  or  Cu++  ion  present  in 
the  saturated  solutions  with  NHs,  forming  the  complex  cation  Cd(NHa)4++  or 
Cu(NHa)4++.  These  complex  cations  have  an  extremely  small  ionization 
tendency;  thus  in  a  normal  NH^OH  solution  the  ratio  of  the  concentration 
of  the  complex  cadmium  ion  to  the  simple  cadmium  ion  is  about  107.  The 
solubility  of  these  hydroxides  in  NH^OH  is  greatly  increased  by  the  presence 
of  ammonium  salts,  since  these  salts,  owing  to  the  common-ion  effect,  greatly 
reduce  the  ionization  of  the  NH4OH,  and  therefore  the  OH~  concentration 
in  the  solution,  thus  enabling  the  Cd++  or  Cu++  concentration,  and  there- 
fore also  the  corresponding  complex  ion  concentration,  to  attain  a  much  larger 
value  than  in  the  saturated  solutions  of  Cd(OH)2  or  Cu(OH)2  in  NBUOH  alone. 

Procedure  36. — Confirmatory  Test  for  Bismuth. — Pour  through  the 
filter  containing  well  washed  NH4OH  precipitate  (P.  35)  a  cold  freshly 
prepared  solution  of  sodium  stannite  (see  Note  1).  (Black  residue, 
presence  of  BISMUTH.) 


58  ANALYSIS  OF  THE  COPPER-GROUP.  P.  86 


Notes.  —  1.  The  solution  of  sodium  stannite  (NaaSnC^)  is  prepared  when 
needed  by  adding  NaOH  solution,  a  few  drops  at  a  time,  to  1  cc.  SnCk  reagent 
diluted  with  5  cc.  water,  until  the  Sn(OH)2  first  formed  is  dissolved  and  a 
clear  or  slightly  turbid  liquid  results.  The  solution  must  be  freshly  prepared, 
because  it  decomposes  spontaneously  into  sodium  stannate  (Na^SnOs)  and 
metallic  tin,  and  because  it  oxidizes  in  contact  with  air  to  sodium  stannate. 
SnC^Bfe  is  an  example  of  a  so-called  amphoteric  substance  —  one  which  acts 
either  as  a  base  or  an  acid,  as  is  shown  by  its  solubility  in  both  acids  and 
alkalies. 

2.  The  final  test  with  sodium  stannite  depends  on  the  reduction  of  Bi(OH)s 
to  black  metallic  bismuth.  The  test  is  an  extremely  delicate  one.  The  other 
reducible  substances,  like  HSb03,  Fe(OH)3,  Pb(OH)2,  or  Cu(OH)2,  that  might 
possibly  be  present  in  the  NH^OH  precipitate  are  not  reduced  by  short  con- 
tact with  stannite  solution  in  the  cold.  Mercury,  if  present,  would  cause 
blackening;  but  it  is  not  precipitated  by  NH^OH  in  the  presence  of  ammo- 
nium salt. 

Procedure  37.  —  Confirmatory  Test  for  Copper.  —  Acidify  one- 
fourth  of  the  NH4OH  solution  (P.  35)  with  HAc,  add  one  drop 
K4Fe(CN)e  solution,  and  allow  the  mixture  to  stand  for  2  or  3 
minutes.  (Red  precipitate,  presence  of  COPPER.)  Then  add  4-5  cc. 
more  K4Fe(CN)6  solution.  If  it  is  uncertain  whether  there  is  a  pre- 
cipitate, pour  the  solution  through  a  filter  and  wash  with  a  little 
water.  (Pink  color  on  the  filter,  presence  of  COPPER.) 

Notes.  —  1.  The  confirmatory  test  for  copper  is  more  delicate  than  the 
formation  of  a  blue  color  with  NHjOH  (P.  35).  It  should,  therefore,  be  tried 
even  when  the  NH^OH  solution  is  colorless.  Cadmium  is  also  precipitated 
by  K4Fe(CN)e;  but  the  precipitate  is  white,  and  does  not  prevent  the  pink 
color  of  the  copper  compound  from  being  detected,  provided  only  a  small 
quantity  of  K4Fe(CN)e  is  added  ;  for  the  copper  salt,  owing  to  its  smaller 
solubility,  is  first  precipitated. 

2.  Nickel,  like  copper,  givs  a  blue  solution  with  excess  of  NH40H;  but 
even  if  present  it  would  giv  a  green,  not  a  red,  precipitate  with  K4Fe(CN)e. 

Procedure  38.  —  Detection  of  Cadmium.  —  Treat  the  remainder 
of  the  NH4OH  solution  (P.  35)  with  KCN  solution  (see  Note  1), 
adding  only  a  few  drops  if  the  solution  is  colorless,  but  enough  to 
decolorize  it  if  it  is  blue,  and  pass  in  H^S  gas  for  about  half  a  minute. 
(Immediate  yellow  precipitate,  presence  of  CADMIUM.) 

Notes.  —  1.  In  working  with  KCN  bear  in  mind  that  it  is  extremely  poisonous. 
Take  care  not  to  get  it  on  the  hands;  also  not  to  breathe  its  fumes,  especially  when 
a  solution  containing  it  is  acidified. 

2.  By  the  addition  of  KCN  the  copper  salt  is  reduced  from  the  cupric  to 
the  cuprous  state  and  then  combines  with  the  excess  of  KCN  to  form  the  com- 
plex salt  K+3Cu(CN)<r  (potassium  cuprocyanide).  This  result  is  due  to  the 
fact  that  cupric  cyanide  tends  to  decompose  spontaneously  into  cuprous  cyanide 
(CuCN)  and  cyanogen  (CN)2,  and  that  this  reaction  takes  place  completely 


P.S8  ANALYSIS  OF  THE  COPPER-GROUP.  59 

in  KCN  solution,  owing  to  removal  of  the  CuCN  by  combination  with  the  excess 
of  KCN.  In  the  presence  of  NH^OH  the  cyanogen  is  not  evolved  as  a  gas, 
but  reacts  with  it,  forming  cyanate  (NEUCNO)  and  cyanide  and  other  more 
complex  products.  The  fact  that  neither  CuS  nor  Cu2§  is  precipitated  from 
this  solution  by  an  alkaline  sulfide  shows  that  neither  the  Cu++  nor  Cu+ 
concentration  is  sufficient  to  cause  the  value  of  the  product  (Cu++)  X  (S~)  or 
(Cu+)2  X  (S=)  to  attain  that  prevailing  in  solutions  saturated  with  CuS  or  Cu2S. 
The  extremely  small  Cu+  concentration  is  due  to  the  very  slight  ionization 
tendency  of  the  complex  anion  Cu(CN)4~;  thus  it  has  been  estimated  that  in  a 
normal  KCN  solution  the  ratio  of  the  concentration  of  the  complex  anion  to 
that  of  the  simple  Cu+  ion  is  about  1026. 

3.  The  cadmium  ammonia  salt  is  also  converted  by  KCN  into  a  complex 
cyanide,  namely,    K+2Cd(CN)4=    (potassium   cadmi cyanide),    since    its   com- 
plex anion  has  a  much  smaller  ionization  tendency  than  the  cation  Cd(NHs)4++; 
thus  in  a  normal  KCN  solution  the  ratio  of  the  concentration  of  the  complex 
anion  to  that  of  the  simple  Cd^+  ion  is  about  1017.    Yet  this  complex  anion  is 
sufficiently   dissociated   into   Cd++  to  cause  the  solution  to  become    super- 
saturated with  CdS  when  an  alkaline  sulfide  is  added. 

4.  The  presence  of  cadmium  is  shown  by  the  immediate  formation  of  a  yellow 
precipitate  by  the  H^S,  but  not  by  a  yellow  coloration  of  the  solution  nor  by 
the  separation  of  a  precipitate  upon  standing.    For,  when  much  copper  is 
present  and  the  solution  is  saturated  with  H^S  and  allowed  to  stand,  the  solu- 
tion soon  becomes  of  a  deep  yellow  color  and  an  orange-red  crystalline  precip- 
itate may  later  separate  from  it,  owing  to  the  fact  that  the  orange-red  com- 
pound (CSNH2)2  is  formed  by  union  of  the  H^S  with  the  (CN)2  set  free  by 
the  reduction  of  the  Cu(CN)2- 

5.  A  very  small  black  precipitate  (which  may  be  due  to  HgS  or  PbS)  may 
sometimes  be  produced  in  the  final  test  for  cadmium  with  E^S  ;  but,  provided 
the  analysis  has  been  properly  conducted,  not  in  sufficient  quantity  to  pre- 
vent the  yellow  color  of  1  or  2  mg.  of  CdS  from  being  seen.    In  case  a  black 
precipitate  is  produced,  and  thus  prevents  a  positiv  conclusion  as  to  the  pres- 
ence or  absence  of  a  little  cadmium,  the  precipitate  may  be  treated,  in  order 
to  eliminate  the  black  sulfide,  as  follows:  Boil  the  precipitate  gently  for  about 
5  minutes  in  a  covered  casserole  with  about  15  cc.  of  1.2-normal  EfeSO^  filter, 
cool  the  filtrate,  add  to  it  three  times  its  volume  of  water,  and  pass  EfeS  into 
it  for  5-10  minutes.    A  yellow  precipitate  of  CdS  will  then  be  obtained,  if  cad- 
mium is  present. 


60 


ANALYSIS  OF  THE  TIN-GROUP. 
ANALYSIS   OF   THE   TIN-GROUP. 


P.  41 


TABLE  VI. — ANALYSIS  OF  THE  TIN-GROUP. 


PRECIPITATE  FROM  AMMONIUM  SULFIDE  SOLUTION : 

Heat  with  10  cc.  12-normal  HCl  (P.  41). 


Solution:  SbCla,  SnCU.     Dilute  to  50  cc.,  heat, 
and  pass  in  HyS  (P.  44). 

Residue:  As2S5. 
Dissolve  in  HCl 
and  KClOs  (P.  #B). 

Orange  precipitate:  80283. 
Dissolve  in  HCl,    add 
Sn  and  Pi  (P.  45). 

Solution:  SnCU. 
Cool,  dilute,  pass  in 
HzS  (P.  46). 

Solution:  HsAsO^ 
Add  NH£)H,  NH*Cl, 
and  MgCh  (P.  48)  . 

Black  deposit:  Sb. 
Treat  with  NaOCl. 

Yellow  precipitate:  81182. 
Evaporate  without  filtering, 
add  Pb,  boil  (P.  46.) 

White  precipitate: 
MgNH4As04. 
Dissolve  in  HCl, 
add  H2S  (P.  43). 

Black  deposit:  Sb. 

Solution:  SnCfe. 
Add  HgCk  (P.  46}. 

Yellow  precipitate: 

AS2S5,   Ai^Ss, 

and  8. 

White  precipitate: 
Hg2Cl2. 

Procedure  41. — Treatment  of  the  Sulfides  with  Strong  Hydrochloric 
Acid. — Transfer  the  precipitated  sulfides  dried  by  suction  (P.  23)  to 
a  test-tube,  add  from  a  small  graduate  exactly  10  cc.  12-normal  HCl, 
place  the  test-tube  in  a  small  beaker  of  water,  heat  the  water  till  the 
contents  of  the  tube  begin  to  boil,  and  then  keep  the  water  for  ten 
minutes  at  a  temperature  which  causes  only  slight  bubbling  in  the 
tube,  stirring  its  contents  from  time  to  time.  Add  3  cc.  water,  and 
filter  with  the  aid  of  suction.  Remove  the  filtrate,  and  treat  it  by 
P.  44.  Wash  the  residue  with  6-normal  HCl,  and  treat  it  by  P.  42. 

Notes. — 1.  If  a  much  weaker  HCl  solution  than  the  acid  of  12-normal  con- 
centration is  used,  or  if  the  acid  becomes  diluted  by  an  unnecessary  quantity  of 
water  left  in  the  precipitate,  much  80285  will  be  left  undissolved.  Even  with 
the  strong  acid  some  8)3285  may  remain  undissolved,  especially  when  a  large 
quantity  is  present,  in  which  case  the  residue  if  small  in  amount  will  have  an 
orange  color.  This  small  quantity  of  80285  would  be  only  very  slowly  removed 
by  further  treatments  with  HCl;  it  does  not,  however,  interfere  with  the  sub- 
sequent tests  for  arsenic.  Moreover,  when  only  a  small  quantity  of  80285  is 
originally  present,  a  large  proportion  of  it  is  extracted,  so  that  it  will  not  escape 
detection.  80285  dissolves  with  formation  of  SbCla,  the  element  being  reduced 


P.  41  ANALYSIS  OF  THE  TIN-GROUP.  61 

from  the  antimonic  state  by  the  K^S  with  liberation  of  sulfur;  81182  dissolves 
with  formation  of  SnCU. 

2.  If  the  solution  be  heated  so  that  only  slight  bubbling  occurs  during  the 
treatment  with  HC1,  the  amount  of  As2S5  which  dissolves  in  ten  minutes  is 
insignificant.     But  this  is  no  longer  true  if  the  solution  be  allowed  to  boil  rapidly; 
for  the  boiling  expels  from  the  solution  the  H^S  liberated  from  the  other  sulfides 
or  by  slight  decomposition  of  the  As2§5  itself,  and  thus  enables  the  decomposi- 
tion of  the  latter  to  proceed  further. 

3.  As2Sa  is  more  rapidly  dissolved  by  HC1  than  is  As2S5.     If  the  former 
can  be  present  in  the  precipitate  (which  can  be  the  case  only  when  ammonium 
monosulfide  was  used  for  separating  the  copper  and  tin  groups),  the  procedure 
should  be  modified  by  saturating  with  E^S  gas  the  cold  concentrated  HC1 
with  which  the  sulfides  are  treated,  and  by  passing  a  slow  current  of  E^S  gas 
through  the  mixture  during  the  heating.    Under  these  conditions  scarcely  any 
As^Sa  dissolves. 

4.  About  3  cc.  water  are  added  to  the  HC1  solution  to  enable  it  to  be  fil- 
tered.   If  more  is  added  and  the  H2S  has  not  all  been  expelled  from  the  solu- 
tion, a  precipitate  of  Sb2Ss  may  separate.    If  this  happens  after  the  filtration, 
it  does,  of  course,  no  harm. 

5.  Care  must  be  taken  to  follow  closely  the  directions  in  regard  to  the 
quantity  of  HC1  used  and  to  avoid  any  loss  of  the  solution  in  the  filtration;  for 
the  subsequent  separation  of  antimony  and  tin  (P.  44)  depends  upon  a  proper 
concentration  of  the  acid. 

6.  The  greater  part  of  any  CuS  and  HgS  present  will  be  dissolved  by  the 
HC1,  and  wiU  be  precipitated  later  with  the  Sb2S3  (P.  44).    A  little  remains 
with  the  As2S5,  but  this  does  not  interfere  with  the  tests  for  arsenic. 

Procedure  42. — Detection  of  Arsenic. — Warm  the  residue  from 
the  HC1  treatment  (P.  41)  with  5-10  cc.  6-normal  HC1,  adding  solid 
KClOa  0.1  g.  at  a  time  until  the  reaction  is  complete;  filter  off  the 
sulfur,  and  evaporate  the  solution  to  about  2  cc.  Add  NH4OH  gradu- 
ally until  the  solution  after  shaking  smells  of  it;  cool,  filter  off  and 
reject  any  precipitate.  Add  to  the  filtrate  in  a  test-tube  about  one- 
third  its  volume  of  15-normal  NH4OH  and  several  drops  of  mag- 
nesium ammonium  chloride  reagent,  and  shake.  If  no  precipitate 
appears,  rub  the  walls  of  the  test-tube  gently  with  a  glass  rod  for  a 
minute  or  two.  (White  crystalline  precipitate,  presence  of  ARSENIC.) 
Collect  the  precipitate  on  a  filter  and  wash  it  once  with  6-normal 
NH4OH.  (Precipitate,  P.  43;  filtrate,  reject.) 

Notes. — 1.  The  main  reaction  between  KClOs  and  concentrated  HC1  is  |bne 
formation  of  Ck;  the  yellow  color  results  from  the  formation  of  a  small  pro- 
portion of  chlorine  dioxide,  CIC^. 

2.  As2S5,  tho  only  very  slowly  dissolved  by  HC1  alone,  is  dissolved  rapidly 
by  it  in  the  presence  of  Ck,  because  of  the  destruction  by  oxidation  of  the 
sulfide-ion  and  of  the  E^S  formed  from  it.  The  same  principles  are  involved 
as  in  the  action  of  HNOs  on  sulfides  (see  P.  31,  Note  2).  It  is  dissolved  with 
formation  of  HsAsC^J  the  proportion  of  AsCl$  that  exists  even  in  a  strong 


62  ANALYSIS  OF  THE  TIN-GROUP.  P.  4* 

HC1  solution  is  extremely  small.  When,  as  here,  arsenic  is  present  in  the 
higher  state  of  oxidation,  solutions  of  it  may  be  boiled  without  loss  of  an 
amount  of  arsenic  significant  in  qualitativ  analysis. 

3.  A  white  precipitate  produced  on  adding  NHiOH  may  arise  from  the 
presence  of  mercury.     The  NKiOH  solution  may  contain  not  only  arsenic, 
but  also  the  small  quantities  of  copper  (if  ammonium  polysulfide  was  used), 
antimony,  and  stannic  tin  that  were  not  dissolved  out  of  the  sulfide  precipitate 
by  HC1. 

4.  The  test  for  arsenic  depends  on  the  formation  of  magnesium  ammonium 
arsenate,  Mg(NH4)AsO4-     This  salt  is  somewhat  soluble  even  in  cold  water, 
and  therefore  the  solution  tested  should  be  fairly  concentrated.     Owing  to 
hydrolysis  (into  NE^OH  and  Mg++HAsO4=),  the  precipitate  is  much  more 
soluble  in  water  than  in  a  strong  NB^OH  solution;   hence  the  addition  of  a 
large  quantity  of  the  latter.     Like  other  crystalline  precipitates,  it  tends  to 
form  a  supersaturated  solution.     Precipitation  is  promoted  by  agitation,  by 
rubbing  the  walls  of  the  tube  with  a  glass  rod,  and  by  increasing  the  degree 
of  supersaturation,  which  is  done  by  concentrating  the  solution  and  adding 
NEUOH.    Provided  these  precautions  are  taken  and  the  total  volume  of  the  final 
solution  does  not  exceed  5  cc.,  the  presence  of  0.5  mg.  of  arsenic  can  be  detected. 
Care  must  be  taken  not  to  scratch  the  glass  by  violent  rubbing,  since  the  pow- 
dered glass  may  be  mistaken  for  the  MgNH4As04  precipitate. 

5.  The  magnesium  ammonium  chloride  reagent  contains  MgCl2  and  NE^Cl. 
The  presence  of  the  latter  salt,  by  reducing  the  OH~  concentration,  prevents 
the  precipitation  of  Mg(OH)2  by  NH4OH. 

Procedure  43. — Confirmatory  Test  for  Arsenic. — Dissolve  the 
MgCl2.NH4Cl  precipitate  (P.  42)  by  pouring  a  5-10  cc.  HC1  through 
the  filter,  saturate  the  solution  with  H2S,  heat  it  nearly  to  boiling, 
and  pass  in  H2S  for  5  minutes.  (White  precipitate  turning  yellow, 
presence  of  ARSENIC.) 

Notes. — 1.  The  slow  formation  of  a  pale  yellow  precipitate  with  H2S  is  a 
characteristic  test  for  HsAsC^.  The  precipitate  is  a  mixture  of  A^Ss,  As^s, 
and  sulfur;  the  proportion  of  As2S5  being  smaller,  the  higher  the  temperature 
at  which  the  precipitation  takes  place  and  the  smaller  the  concentration  of 
the  HC1. 

2.  A  considerable  amount  of  H2S  is  absorbed  by  a  cold  HsAsC^  solution 
before  any  precipitate  appears.  This  is  due  to  the  conversion  of  a  part  of  the 
HgAsC)4  into  HsAsOsS,  which  then  decomposes  slowly,  giving  HsAsO3  and 
sulfur.  This  last  decomposition  is  accelerated  by  increasing  the  H+  con- 
centration and  by  raising  the  temperature.  The  HsAsOa  formed  reacts  at 
once  with  H2S,  and  A^Ss  is  precipitated.  The  reactions  taking  place  are: 

H2S  =  H3As03S+H20, 

=H3AsO3  +S, 

2H3AsO3+3H2S  =  As2S3       +6H2O, 
which  together  giv  the  resultant  reaction, 

2HsAsO4+5H2S  =  As2S3+S2+8H2O. 

The  rate  of  this  reaction  depends  upon  the  rate  of  the  slowest  of  the  separate 
reactions — the  decomposition  of  the  HsAsO3S — and  on  the  quantity  of  this 


P.  43  ANALYSIS  OF  THE  TIN-GROUP.  63 


substance  which  has  been  produced  by  the  first  reaction.  This  quantity 
to  be  determined  by  the  equilibrium-conditions  of  the  first  reaction  rather 
than  by  its  rate;  for  it  is  larger  when  the  solution  is  saturated  in  the  cold  with 
H2S,  doubtless  owing  to  the  greater  solubility  of  the  H2S  hi  the  cold  solution. 
This  explains  why  it  is  advantageous  to  saturate  the  solution  with  H2S  first 
in  the  cold.  Having  formed  in  this  way  as  large  a  proportion  of  HsAsOaS  as 
possible,  the  solution  is  then  heated  in  order  to  cause  decomposition  of  this 
substance  according  to  the  second  reaction. 
3 .  As-jSs  results  from  an  independent  reaction,  taking  place  slowly,  as  follows : 

2H3As04+5H2S = As2S5+4H20. 

The  fact  that  it  is  the  main  product  of  the  action  of  the  H2S  when  the  solu- 
tion is  cold  and  concentrated  in  HC1  is  doubtless  due  to  the  conversion  of  more 
of  the  HsAsO4  into  AsCls  under  these  conditions  (which  are  those  least  favor- 
able to  hydrolysis)  and  to  the  partial  ionization  of  the  AsCls  into  As+++++ 
and  Cl~  ions. 

Procedure  44. — Separation  of  Antimony  and  Tin. — Dilute  the  solu- 
tion from  the  HC1  treatment  of  the  sulfides  (P.  41)  with  water  to  a  vol- 
ume of  55  cc.,  transfer  it  to  a  small  flask  placed  in  a  400-cc.  beaker  of 
water,  heat  the  water  to  boiling,  and  pass  into  the  solution  a  moderate 
current  of  H2S  gas  for  8-10  minutes  but  not  longer,  keeping  the  water 
in  the  beaker  gently  boiling.  (Orange-red  precipitate,  presence  of 
ANTIMONY.)  Filter  while  hot,  and  wash  the  precipitate  with  hot  water. 
(Precipitate,  P.  45;  filtrate,  P.  46.) 

Notes. — 1.  By  following  carefully  the  directions  given  in  P.  41  and  in  this 
procedure,  a  good  separation  of  antimony  and  tin  may  be  obtained;  thus, 
when  only  1  mg.  of  antimony  is  present  it  is  precipitated,  while  even  500  mg. 
of  (stannic)  tin  giv  no  precipitate.  If,  however,  the  HC1  solution  be  more  con- 
centrated, a  small  quantity  of  antimony  will  escape  detection.  On  the  other 
hand,  if  the  HC1  solution  be  more  dilute,  or  if  it  be  not  kept  hot,  some  SnS2 
may  precipitate  when  a  large  amount  of  tin  is  present.  When  SnS2  is  mixed 
with  a  little  Sb2S3  a  brown  precipitate  results. 

2.  If  mercury  or  copper  be  present  in  the  substance  and  ammonium  poly- 
sulfide  has  been  used  in  separating  the  copper  and  tin  groups,  HgS  or  CuS  may  be 
precipitated  at  this  point  as  a  gray  or  black  precipitate. 

Procedure  45. — Confirmatory  Test  for  Antimony. — Dissolve  the 
H2S  precipitate  (P.  44)  in  a  little  12-normal  HC1  in  a  small  casserole, 
and  evaporate  the  solution  to  about  1  cc.  Cool,  introduce  beneath 
the  solution  a  strip  of  platinum  foil,  and  place  upon  the  platinum  a 
layer  of  granulated  tin.  After  several  minutes  wash  the  platinum  foil 
carefully  with  water,  and  immerse  it  in  NaOCl  solution.  (Black  deposit 
on  the  platinum  undissolved  by  NaOCl,  presence  of  ANTIMONY.) 

Notes. — 1.  Mercury  and  copper,  if  present,  will  also  be  precipitated  in  the 
metallic  condition  upon  the  platinum,  but  the  antimony  may  be  easily  distin- 
guished from  the  gray  deposit  of  mercury  or  the  reddish  one  of  copper  by  its 


64  ANALYSIS  OF  THE  TIN-GROUP.  P.  45 

coal-black  color.     Tin  is  used,  rather  than  zinc,  in  precipitating  the  antimony, 
since  zinc  would  also  precipitate  tin  from  the  solution. 

2.  By  the  contact  of  the  platinum  with  the  tin  a  voltaic  cell  is  formed, 
with  the  result  that  the  hydrogen  is  liberated  on  the  surface  of  the  platinum, 
instead  of  on  that  of  the  tin,  as  it  would  be  if  tin  alone  were  used.     This  has 
the  advantages  of  accelerating  the  action  of  the  acid  on  the  tin,  and  of  causing 
the  antimony  to  deposit  on  the  platinum,  where  it  can  be  more  readily  seen. 

3.  The  treatment  with  NaOCl  serves  to  prove  that  the  black  precipitate 
does  not  consist  of  arsenic  ;   for  this  element  is  readily  dissolved  by  it,  while 
antimony,  copper,  or  mercury  is  not.    Since,  however,  5-10  mg.  of  arsenic  must 
be  present  before  the  treatment  with  tin  would  giv  a  deposit  on  the  platinum, 
an  arsenic  deposit  is  not  likely  to  be  obtained  in  a  properly  conducted  analysis. 

4.  The  quantity  of  antimony  present  can  usually  be  better  estimated  from  the 
size  of  the  H2S  precipitate  than  from  the  appearance  of  the  black  deposit  on  the 
platinum. 

Procedure  46. — Detection  of  Tin. — Cool  the  filtrate  from  the  H2S 
precipitate  (P.  44),  dilute  it  with  25  cc.  water,  saturate  it  with  H2S, 
cork  the  flask,  and  let  it  stand  for  10  minutes.  (Yellow  precipitate, 
presence  of  TIN.) 

Confirmatory  Test  for  Tin. — If  H2S  has  produced  a  precipitate, 
evaporate  the  mixture  without  filtering  to  2-3  cc.,  pour  it  into  a  small 
conical  flask,  add  10  cc.  water  and  10  g.  of  finely  granulated  lead,  cover 
the  flask  with  a  small  watch-glass,  and  boil  slowly  for  5  to  10  minutes. 
Pour  the  hot  solution  through  a  filter  into  10  cc.  0.2  normal  HgCl2  solu- 
tion. If  there  is  a  precipitate,  heat  the  mixture  to  boiling.  (White 
precipitate,  presence  of  TIN.) 

Notes. — 1.  The  solution  is  precipitated  with  H2S  in  the  cold,  because  a 
small  quantity  of  SnS2  would  not  separate  from  a  hot  solution  unless  the  acid 
were  more  diluted.  The  addition  of  much  water  is  avoided,  since  it  has  to  be 
evaporated  off  in  the  confirmatory  test.  When  only  a  small  quantity  (0.5-2 
mg.)  of  tin  is  present,  the  precipitate  of  SnS2  produces  in  the  solution  a  yellow- 
ish translucent  turbidity,  which  is  readily  distinguishable  from  the  trace  of 
finely  divided  sulfur  which  may  separate. 

2.  With  respect  to  the  confirmatory  test  the  following  points  deserve  men- 
tion. The  precipitate  of  SnS2  is  not  filtered  off  but  is  dissolved  by  concentrating 
the  acid  by  evaporation,  since  it  clogs  the  filter  and  tends  to  pass  through  it. 
Since  SnCl2  oxidizes  rapidly  in  the  air,  theksolution  in  HC1  is  immediately  added 
to  the  HgClj  solution.  The  mixtures  fnmlly  heated  to  boiling  to  make  sure 
that  the  ^precipitate  iS  not  PbClj.  •«.?  f- 


P.  51      PRECIPITATION  OF  ALUMINUM  AND  IRON  GROUPS. 


65 


PRECIPITATION  AND  SEPARATION  OF  THE  ALUMINUM  AND  IRON  GROUPS. 


TABLE  VII. — PRECIPITATION  AND  SEPARATION  OF  THE  ALUMINUM  AND  IRON  GROUPS. 


FILTRATE  FROM  THE  HYDROGEN  SULFIDE  PRECIPITATE. 

Add  NH£)H  in  excess  (P.  51}. 


Precipitate*:  A1(OH)8,  Cr(OH)3,  Fe(OH)2_3. 

Solution:  Salts  of  ZnCNHsk  Co(NH3)4,  Ni(NH3)4,  Mn,  Ba,  Sr,  Ca,  Mg,  K,  and  Na. 
Add  (NHJzS  and  filter  (P.  51}. 


Precipitate*:  A1(OH)3,  Cr(OH)3,  FeS,  ZnS,  MnS,  CoS,  NiS. 
Dissolve  in  HCl  and  HN03,  add  NaOH  (P.  52}. 


Precipitate*:  Fe(OH)3,  Mn(OH)2,  Co(OH)2,  Ni(OH)2. 
Solution:   NaAlO2,  NaCrO2,  Na2Zn02. 
Add  Na20z  and  filter  (P.  52}. 


Filtrate:     NaAlO2,    Na2Cr04, 
Na2ZnO2. 
See  Table  VIII. 


Precipitate*:  Fe(OH)3, 
MnO(OH)2,  Co(OH)3, 
Ni(OH)2. 

See  Table  IX. 


Filtrate:    Salts  of  Ba, 
Sr,  Ca,  Mg,  K,  Na. 


*  When  phosphate  is  present  in  the  solution,  these  precipitates  may  contain  the  phosphates  of  the 
elements  otherwise  precipitated  as  hydroxides,  and  also  the  phosphates  of  barium,  strontium,  calcium, 
and  magnesium. 

Procedure  51. — Precipitation  of  the  Aluminum  and  Iron  Groups. — 
Boil  the  filtrate  from  the  H2S  precipitate  (P.  21)  till  the  H2S  is  ex- 
pelled. Add  to  it  10-15  cc.  NH4OH,  shake,  and  note  whether  there  is 
a  precipitate.  Add  ammonium  monosulfide  slowly  (or,  in  case  nickel 
seems  to  be  present,  pass  in  H2S)  until,  after  shaking,  the  vapors 
in  the  flask  blacken  moist  PbAc2  paper.  Heat  the  mixture  nearly  to 
boiling,  shake  it,  and  let  it  stand  2  or  3  minutes.  (Precipitate,  pre- 
sence Of  ALUMINUM-GROUP  OT  IRON-GROUP  Or  of  ALKALINE-EARTH 

PHOSPHATE.)  Filter,  and  wash  the  precipitate,  first  with  water  con- 
taining about  1%  of  the  (NH4)2S  reagent,  and  then  with  a  little  pure 
water.  If  the  filtration  is  slow,  keep  the  funnel  covered  with  a  watch- 
glass.  To  the  filtrate  add  a  few  drops  (NH4)2S,  boil  the  mixture  for 
a  few  seconds  (or,  in  case  it  is  dark  colored,  until  it  becomes  colorless 
or  light  yellow);  filter  if  there  is  a  precipitate,  uniting  it  with  the 
preceding  one.  (Precipitate,  P.  52;  filtrate,  P.  81.) 


66          PRECIPITATION  OF  ALUMINUM  AND  IRON  GROUPS.      P.  61 

Notes. — 1.  The  H2S  is  boiled  out  and  the  effect  of  the  addition  of  NH4OH 
alone  is  noted  because  it  often  givs  a  useful  indication  as  to  what  elements  are 
present.  To  save  time  the  expulsion  of  the  £[28  may  be  omitted  when  this 
indication  is  considered  unimportant.  Only  a  slight  excess  of  ammonium 
monosulfide  is  used,  in  order  to  prevent  as  far  as  possible  dissolving  the  NiS. 
By  passing  £[28  into  the  ammoniacal  solution,  instead  of  adding  (NEL^S,  the 
dissolving  of  NiS  is  entirely  prevented;  therefore,  tho  the  operation  takes  a 
little  longer,  the  use  of  H2S  is  to  be  preferred  when  nickel  is  likely  to  be  present. 
The  mixture  is  shaken  in  order  to  coagulate  the  precipitate  and  make  it  filter 
more  readily.  Heating  also  promotes  the  coagulation  of  the  precipitate:  heat 
is  therefore  applied  when  the  precipitate  does  not  coagulate  and  settle  quickly 
on  shaking.  The  filtrate  is  boiled  for  a  few  moments  to  ensure  the  complete 
precipitation  of  Cr(OH)s,  or  longer  to  ensure  that  of  NiS,  whose  presence  is 
indicated  by  a  brown  or  nearly  black  color  of  the  filtrate.  Finally,  it  is  directed 
to  wash  with  water  containing  a  little  (NKj^S  and  to  keep  the  filter  covered, 
in  order  that  some  excess  of  (NH^S  may  always  be  present ;  for,  if  the  (NH^S 
adhering  to  the  precipitate  is  removed  by  oxidation  or  by  volatilization  (as 
H2S  and  NHs),  the  sulfides  are  oxidized  to  soluble  sulfates  by  the  air. 

2.  Under  the  conditions  of  the  procedure,  which  provides  for  a  small  excess 
of  NB^OH  in  the  presence  of  ammonium  salt,  aluminum,  chromium,  and  iron 
are  completely  precipitated.    Al(OH)s  is  white;   Cr(OH)3,  grayish  green.    The 
color  of  the  precipitated  iron  hydroxide  varies  with  the  state  of  oxidation  of  the 
iron,  pure  ferrous  salts  yielding  a  white  precipitate,  and  ferric  salts  a  reddish 
brown  one,  while  mixtures  of  them  yield  green  or  black  precipitates  ;  in  the 
alkaline  mixture  the  precipitate  is  rapidly  oxidized  by  the  oxygen  of  the  air  and 
undergoes  corresponding  changes  in  color.    Under  the  conditions  of  the  proce- 
dure manganese  is  not  precipitated  as  Mn(OH)2j  but  in  the  alkaline  solution 
the  manganous  salt  is  rapidly  oxidized  by  the  air  with  the  formation  of  a 
brown  precipitate  consisting  of  Mn(OH)s  and  MnOaE^  (a  hydrate  of  MnO2). 
Zinc,  nickel,  and  the  alkaline-earth  elements  remain  in  solution;  nickel  yield- 
ing a  blue  solution.    Cobalt  also  remains  dissolved  (yielding  a  pink  solution), 
unless  the  quantity  present  is  large,  in  which  case  a  blue  precipitate  may 
separate;  owing  to  oxidation  the  solution  rapidly  changes  to  a  dark  orange, 
and  the  precipitate  to  a  bright  green  color.     When  much  chromium  is  present, 
zinc  and  magnesium  may  be  completely  precipitated  in  combination  with  it. 
If  a  much  larger  excess  of  NEUOH  is  employed  than  is  directed,  a  few  milli- 
grams of  aluminum  and  chromium  may  be  dissolved,  the  latter  giving  a  pink 
colored  solution. 

3.  The  presence  of  ammonium  salts  in  the  solution  serves  to  prevent  the 
precipitation  of  Mg(OH)2,  and  also  to  lessen  the  amount  of  Al(OH)s  dissolved 
by  the  NEUOH. 

4.  The  influence  of  an  excess  of  the  NI^OH  and  of  the  presence  of  am- 
monium salt  on  the  solubilities  of  the  various  hydroxides  is  explained  by  the 
mass-action  law  and  ionic  theory  as  follows:    In  order  that  any  hydroxide, 
say  of  the  type  MOoE^,  may  be  precipitated,  it  is  necessary  that  the  product 
(M++)  X  (OH~)2  of  the  concentrations  of  the  ions  M++  and  OH~  in  the  solu- 
tion under  consideration  attain  the  saturation-value  of  that  product.    This 
saturation- value  varies,  of  course,  with  the  nature  of  the  hydroxide;  but  for 
all  the  elements  of  the  iron-group  and  for  magnesium,  it  is  so  small  that 


P.  61      PRECIPITATION  OF  ALUMINUM  AND  IRON  GROUPS.          67 

even  in  a  solution  containing  in  50  cc.  only  1  mg.  of  the  element  and  a  slight 
excess  of  NH4OH,  the  product  (M++)  X  (OH~)2  exceeds  it,  and  precipitation 
results.  When,  however,  much  ammonium  salt  is  also  present,  this  greatly 
reduces,  in  virtue  of  the  common-ion  effect,  the  ionization  of  the  NE^OH  and 
therefore  the  OH~  concentration  in  the  solution,  so  that  now  for  certain  ele- 
ments the  product  (M++)  X  (OH~)2  does  not  reach  the  saturation- value,  even 
when  (M++)  is  moderately  large  (say  500  mg.  in  50  cc.).  This  is  true  of  mag- 
nesium and  manganese;  but  in  the  cases  of  the  trivalent  elements,  aluminum, 
chromium,  and  ferric  iron,  the  solubility  of  the  hydroxides  in  water  is  so  slight 
that  even  in  ammonium  salt  solution  the  solubility  is  not  appreciable. 

If  these  were  the  only  effects  involved,  the  greater  the  excess  of  NB^OH 
added,  the  less  would  be  the  solubility  of  any  hydroxide;  but  other  influences 
come  into  play  with  certain  of  the  elements.  These  influences  are  of  two  kinds. 
The  first  of  these  is  shown  by  zinc,  nickel,  and  cobalt.  In  the  case  of  these 
elements,  just  as  with  silver  and  copper,  the  excess  of  ammonia  combines  with 
the  simple  cation  M++,  forming  complex  cations  of  the  types  M(NH3)2++ 
and  M(NH3)4++,  thereby  removing  the  simple  cation  from  the  solution  and 
making  it  necessary  for  more  of  the  hydroxide  to  dissolve  in  order  to  bring 
back  the  value  of  (M++)  X  (OH~)2  to  the  saturation -value.  In  such  a  case  the 
presence  of  ammonium  salt  increases  the  solubility  still  further  since  it  greatly 
decreases  the  value  of  (OH~),  owing  to  the  common-ion  effect  on  the  ionization 
of  the  NE^OH.  Chromium  also  forms  similar  ammonia  complexes,  but  in 
much  smaller  proportion. 

The  second  kind  of  effect  is  exhibited  in  the  case  of  AlOsHs.  This  hydroxide 
is  another  example  of  an  amphoteric  substance:  for  it  behaves  both  as  a  base 
and  as  an  acid  in  consequence  of  its  being  appreciably  ionized  both  into  OH~ 
and  A1+++  and  into  H+  and  A1O3H2-  (or  into  H+,  AlQr,  and  H2O). 
With  the  H+  arising  from  the  latter  form  of  ionization  the  OH~  coming  from 
the  excess  of  NEUOH  combines  to  form  H^O,  so  as  to  satisfy  the  mass- action 
expression  for  the  ionization  of  water,  (H+)  X  (OH~)  =  a  constant  (which 
has  the  very  small  value  10~14  at  25°).  This  causes  more  AlOsHg  to  dis- 
solve until  the  product  (AKV")  X  (H+)  again  attains  its  saturation-value.  This 
shows  that  the  quantity  of  aluminum  dissolved  increases  with  the  OH~  con- 
centration in  the  solution,  and  that  therefore  it  would  be  much  greater  in  a 
solution  of  a  largely  ionized  base  like  NaOH  than  in  that  of  a  slightly  ionized 
base  like  NELjOH.  It  also  shows  that  the  presence  of  ammonium  salts  tends 
to  neutralize  the  solvent  action  of  an  excess  of  NHiOH,  since  they  decrease  the 
OH~  concentration  in  its  solution. 

5.  It  follows  from  the  statements  in  the  preceding  notes  that,  if  NHiOH 
produces  no  precipitate,  it  proves  the  absence  of  as  much  as  1  mg.  of  aluminum 
and  iron;  also  of  chromium,  if  the  mixture  is  heated  to  boiling  after  the  addi- 
tion of  NKiOH.    Care  must  be  taken  not  to  overlook  a  small  precipitate  which 
might  otherwise  escape  detection  on  account  of  its  transparency.    The  mixture 
should  therefore  be  heated  and  shaken  and  allowed  to  stand  2  or  3  minutes,  in 
order  that  the  precipitate  may  collect  in  flocks.     This  treatment  also  oxidizes 
the  iron  when  present  in  small  quantity,  and  thus  enables  it  to  be  more 
readily  detected;  for  its  precipitation  in  the  ferric  state  is  more  complete. 

6.  When  phosphate  is  present,  magnesium,  calcium,  strontium,  barium, 
and  manganese  may  be  partially,  or  even  completely,  precipitated  by  NH4OH. 


68  SEPARATION  OF  ALUMINUM  AND  IRON  GROUPS.         P.  51 

The  reasons  for  this  are  as  follows.  The  normal  phosphates  and  the  mono- 
hydrogen  phosphates  of  these  elements  are  difficultly  soluble  in  water,  but  dis- 
solve readily  in  acids,  owing  to  the  formation  in  solution  of  the  much  more 
soluble  dihydrogen  phosphates  and  of  free  phosphoric  acid.  Upon  the  addi- 
tion of  an  excess  of  NH^OH  to  such  a  solution  these  acid  compounds  are  con- 
verted into  the  normal  phosphates,  and  these  are  reprecipitated.  It  is  therefore 
necessary,  when  phosphate  is  present,  to  provide  for  the  detection  of  the  alkaline- 
earth  elements  in  the  analysis  of  the  NHjGH  precipitate.  They  are,  however, 
not  necessarily  found  in  that  precipitate;  for,  when  other  elements,  like  iron 
and  aluminum,  which  form  much  less  soluble  phosphates,  are  also  present, 
these  may  combine  with  all  the  phosphate-ion  present,  thus  leaving  the 
alkaline-earth  elements  in  solution. 

7.  The  presence  of  any  other   acidic  constituent  which  forms  with  the 
alkaline-earth  elements  salts  soluble  in  dilute  acids,  but  insoluble  in  ammonia, 
may  also  cause  then*  precipitation  at  this  point.    Fluoride  is  the  only  common 
inorganic  constituent  of  this  kind,  and  it  will  ordinarily  have  been  removed 
in  the  evaporation  with  acids  in  the  preparation  of  the  solution. 

8.  (NH4)2S  precipitates  ZnS,  MnS,  NiS,  and  CoS,  and  converts  Fe(OH)2 
into  FeS,  and  Fe(OH)3  into  F^Sa.    The  hydroxides  of  aluminum  and  chromium 

£'•  are  not  affected  by  the  (NH^S. 

9.  The  sulfides  of  iron,  nickel,  and  cobalt  are  black;    ZnS  is  white;    and 
MnS  is  flesh-colored,  but  turns  brown  on  standing  in  the  air,  owing  to  oxida- 
tion to  Mn(OH)s  and  MnOaH^. 

10.  When  nickel  is  present  alone  or  when  it  forms  a  large  proportion  of 
the  (NH^S  precipitate,  several  milligrams  of  it  usually  pass  into  the  filtrate, 
giving  it  a  brown  or  black  color;    and  some  NiS  also  passes  through  the  filter 
with  the  wash-water.    In  this  case  it  is  useless  to  try  to  remove  the  NiS  by  fil- 
tering again,  but  it  can  be  coagulated  by  boiling  for  several  minutes.     The 
brown  solution  is  formed  only  in  the  presence  of  ammonium  polysulfide.    Its 
formation  can,  as  stated  above,  be  avoided  altogether  by  passing  E^S  into  the 
NHjOH  solution,  instead  of  adding  the  ammonium  monosulfide  reagent,  which 
after  exposure  to  the  air  always  contains  some  polysulfide.    The  nature  of  the 
brown  solution  is  not  known. 

Procedure  52. — Separation  of  the  Aluminum-Group  from  the  Iron- 
Group. — Transfer  the  (NH4)2S  precipitate  (P.  51)-,  wilh  the  filter 
if 'necessary,  to  a  casserole;  add  5-20  cc.  HC1,  stir  for  a  minute  or 
two  in  the  cold,  and  then  boil  the  mixture  for  2  or  3  minutes;  if  a 
black  residue  still  remains,  add  a  few  drops  of  HNOs  and  boil  again. 
Add  5-10  cc.  water,  filter  off  the  sulfur  residue,  and  evaporate  the  fil- 
trate almost  to  dryness  to  remove  the  excess  of  acid.  (See  Note  2, 
P.  12.) 

Dilute  the  solution  to  10  or  20  cc.,  and  make  it  alkaline  with  NaOH 
solution,  avoiding  a  great  excess  and  adding  10-20  cc.  more  water  if 
so  large  a  precipitate  separates  that  the  mixture  becomes  thick  with 
it.  v:  Place  the  casserole  in  a  dish  of  cold  water,  and  add  0.5-2.0  cc. 
powder  in  small  portions  with  constant  stirring.  (See  Note  5.) 


P.  62        SEPARATION  OF  ALUMINUM  AND  IRON  GROUPS.  69 

Then  add  5  cc.  3-normal  Na2C03  solution;  boil  for  2  or  3  minutes  to 
decompose  the  excess  of  Na202,  cool,  dilute  with  an  equal  volume  of 
water,  filter  with  the  help  of  suction,  and  wash  with  hot  water. 
(Precipitate,  P.  61;  filtrate,  P.  53.) 

Notes. — 1.  All  the  hydroxides  and  all  the  sulfides,  except  NiS  and  CoS, 
usually  dissolve  readily  in  cold  HC1.  If,  therefore,  there  is  considerable  black 
residue  after  adding  the  HC1,  it  shows  the  presence  of  nickel  or  cobalt;  a  very 
small  black  residue  may,  however,  be  due  to  FeS  enclosed  within  sulfur.  The 
fact  that  there  is  no  such  dark-colored  residue  does  not,  however,  prove  that 
nickel  and  cobalt  are  entirely  absent;  for  a  considerable  quantity  of  them 
(even  5  mg.)  may  dissolve  completely  in  the  HC1  when  large  quantities  of  other 
elements,  especially  iron,  are  also  present. 

2.  The  (NHOsjS  precipitate  is  first  treated  with  HC1,  partly  in  order  to 
furnish  the  indication  just  referred  to  of  the  presence  of  nickel  or  cobalt,  but 
also  because  much  more  free  sulfur  and  sulfate  would  be  formed  by  oxidation 
if  HNOa  or  aqua  regia  were  used  at  the  start.     (The  presence  of  much  sulfate 
in  the  solution  interferes  with  the  subsequent  test  for  chromate.)     If  NiS  or 
CoS  is  present  in  the  residue,  HNOs  must,  however,  be  subsequently  added, 
to  ensure  the  solution  of  these  sulfides. 

3.  By  NaOH,  iron,  manganese,  nickel,  and  cobalt  are  completely  precipi- 
tated and  do  not  dissolve  in  moderate  excess ;  while  aluminum,  chromium, 
and  zinc  remain  in  solution  or  dissolve  when  a  sufficient  excess  is  added.    The 
solubility  of  the  last  three  elements  is  due  to  the  fact  that  their  hydroxides 
are  amphoteric  substances  which  form    with  the  NaOH  soluble  v  aluminate 
(NaA102),   chromite  (NaCrO2),   and  zincate  (Na^ZnC^),  respectivly.     When 
zinc  and  chromium  are  simultaneously  present  they  are  precipitated  in  the 
form  of  a  double  compound  (ZnCr2O4).    Chromium  would  also  be  completely 
precipitated,  owing  to  hydrolysis  of  the  chromite  and  the  formation  of  a  less 
soluble  solid  hydroxide,  if  the  NaOH  solution  were  boiled  before  adding  Na2Oz. 
Mn(OH)2  is  white,  but  rapidly  turns  brown,  owing  to  oxidation  to  Mn(OH)s; 
Ni(OH)2  is  light  green;  Co(OH)2  is  pink,  but  from  cold  cobaltous  salt  solutions  a 
blue  basic  salt  is  first  precipitated.    If  a  large  excess  of  NaOH  be  added,  a 
little  Co(OH)2  dissolves,  yielding  a  blue  solution,  doubtless  forming  a  salt 
such  as  Na2CoO2.     This  is  to  be  avoided,  since  then  the  cobalt  will  not   be 
completely  oxidized  and  precipitated  upon  the  subsequent  addition  of  Na^. 

4.  By  the  addition  of  Na^,  Fe(OH)2  is  changed  to  dark  red  Fe(OH)3, 
Mn(OH)2  to  brown  hydrated  MnO2,  and  Co(OH)2  to  black  Co(OH)3,  all  of 
which  are  insoluble  in  excess  of  NaOH.    Chromium,  which  after  the  addition 
of  cold  NaOH  is  present  as  soluble  sodium  chromite  (NaCrO2) ,  is  converted  by 
NaA  into  chromate  (Na2CrO4).     This  remains  in  solution  together  with  the 
zinc,  which  is  still  present  as  zincate. 

5.  Even  a  cold  solution  of  Na2O2  decomposes  rapidly  with  evolution  of 
oxygen,  and  this  decomposition  takes  place  with  explosiv  violence  when  the 
solution  is  hot.    The  peroxide  is  therefore  added  in  small  portions  to  the  cold 
solution.     It  is  best  to  transfer  a  little  of  the  powder  from  the  can  in  which  it 
comes  in  trade  to  a  dry  7-cm.  test-tube,  and  then  to  sprinkle  it  slowly  into  the 
solution  with  constant  stirring.     A  steady  evolution  of    gas  continuing  after 
the  mixture  has  been  well  stirred  is  an  indication  that  sufficient  peroxide  has 


70  SEPARATION  OF  ALUMINUM  AND  IRON  GROUPS.  P.  62 

been  added.  When  much  chromium  is  present,  it  should  be  added  till  the  green 
precipitate  disappears  and  the  liquid  assumes  a  dark  yellow  color.  The  solu- 
tion is  diluted  before  filtering  in  order  to  avoid  the  disintegration  of  the  filter- 
paper.  It  is  also  often  advantageous  to  support  the  filter  by  folding  it  together 
with  a  small  hardened  filter. 

6.  This  separation  with  NaOH,  Na^C^,  and  NasCOs  is  a  very  satisfactory 
one,  except  in  the  case  of  zinc.    As  much  as  5  mg.  of  this  element  is  almost 
completely  carried  down  in  the  precipitate  when  much  iron,  nickel,  or  cobalt  is 
present;  and  as  much  as  20  mg.  of  it  may  be  completely  precipitated  when 
much  manganese  is  present.     Provision  is  therefore  made  for  the  detection  of 
zinc  in  the  precipitate. 

7.  The  Na2COs  is  added  to  ensure  the  complete  precipitation  of  magne- 
sium, calcium,  strontium,  and  barium,  whose  hydroxides,  especially  that  of 
barium,  are  somewhat  soluble  even  in  the  presence  of  NaOH.     ZnCOs,  tho 
insoluble  in  a  dilute  solution  of  Na^COa  alone,  dissolves  when  much  NaOH 
is  present,  owing  to  nearly  complete  conversion  of  the  zinc-ion  into  zincate-ion 
(ZnO2=).  The  Na2COs  also  serves  to  decompose  the  chromates  of  the  alkaline-earth 
elements;    if  it  is  not  added,  chromium  may  remain  in  the  precipitate  and 
escape  detection.     It  is  unnecessary  to  add  the  Na^COs  when  the  alkaline- 
earth  elements  are  known  to  be  absent. 

8.  Phosphate,  if  present,  divides  itself  in  this  procedure  between  the  pre- 
cipitate and  solution  in  a  proportion  which  depends  on  the  nature  and  quantities 
of  the  basic  elements  present.     (See  P.  51,  Note  6.)     Its  presence  does  not 
cause  any  of  the  elements  to  precipitate  which  would  not  otherwise  do  so,  in 
spite  of  the  slight  solubility  of  aluminum  and  zinc  phosphates.    This  is  due 
to  the  fact  that  the  cations  of  these  elements  (Al+++,  Zn++)  are  present  in 
the  NaOH  solution  only  at  an  extremely  small  concentration,  owing  to  their 
conversion  by  the  OH~  into  anions  (AIO2~,  ZnO2=). 

9.  If  Na2O2  is  not  available,  sodium  hypobromite,  NaBrO,  may  be  used  as  the 
oxidizing  agent;   but  it  is  not  quite  so  satisfactory  as  Na£O2,  for  it  does  not 
oxidize  Cr(OH)s  so  readily,  and  it  is  apt  to  oxidize  some  of  the  manganese  to 
NaMnO4  (especially  if  there  is  not  a  large  excess  of  NaOH  present). 


P.  5S 


ANALYSIS  OF  THE  ALUMINUM-GROUP. 


71 


ANALYSIS  OP  THE  ALUMINUM-GROUP. 


TABLE  VIII. — ANALYSIS  OP  THE  ALUMINUM-GROUP. 


FILTRATE  FROM  THE  SODIUM  HYDROXIDE  AND    PEROXIDE   TREATMENT! 


NaAlQj, 


Acidify  with  HNOS,  add  NH^OH  (P.  53). 


Precipitate:  A1(OH)3. 
Dissolve  in  HNO*,  add 
Co(NOz)%,  evaporate, 
ignite  (P.  64). 

Filtrate.    Add  HAc  and  BaClz  (P.  55). 

Precipitate:  BaCrO4. 
Dissolve  in  HCl  and 
H2SOz,  evaporate  (P.  66). 

Filtrate:  Zinc  salt. 
Pass  in  H2S  (P.  57). 

Blue  residue: 
Co(A102)2. 

White  precipitate:  ZnS. 
Dissolve  in  HN03,  add 
Co(NOz)z  and  NazCOs, 
ignite  (P.  67). 

Green  color: 
CrCla. 

Green  residue: 
CoZnO2. 

Procedure  53. — Separation  of  Aluminum  from  Chromium  and 
Zinc. — Acidify  the  alkaline  solution  (P.  52)  with  16-normal  HN03, 
avoiding  a  large  excess;  add  NH4OH  until  the  mixture  after  shaking 
smells  of  it,  and  then  add  2-3  cc.  more.  (White  flocculent  precipitate, 
presence  of  ALUMINUM.)  Heat  almost  to  boiling  in  order  to  coagulate 
the  precipitate,  filter,  and  wash  thoroly  with  hot  water.  (Precipitate, 
P.  54;  filtrate,  P.  55.) 

Notes. — 1.  The  alkaline  solution  is  acidified  with  HNOs,  instead  of  with 
HCl,  because  the  latter  acid  might  reduce  chromic  acid,  especially  if  a  large 
quantity  were  added,  or  if  the  acid  solution  were  heated  A  moderate  excess 
of  NH,OH  must  be  added  in  order  to  keep  the  zinc  in  solution,  which  it  does 
because  of  the  production  of  the  complex  cation  Zn(NHs)4++;  but  a  large 
excess  is  to  be  avoided,  since  it  di  solves  A1(OH)3,  owing  to  formation  of 
NH4+A1O2~.  The  zinc  dissolves  even  when  carbonate  or  phosphate  is  present. 

2.  Since  aluminum  and  silica  are  very  likely  to  be  present  in  the  NaOH 
and  Na2O2  used  as  reagents,  a  blank  test  for  these  impurities  should  be  made 
whenever  new  reagents  are  employed  for  the  first  time,  by  treating  20  cc.  of 
the  NaOH  reagent,  or  2  g.  of  NasCfe  added  to  20  cc.  water,  by  P.  53  and  com- 
paring the  NKiOH  precipitate  with  that  obtained  in  the  actual  analysis.  It 
is  also  well  at  the  same  time  to  test  for  zinc  by  acidifying  the  NEUOH  solution 
with  acetic  acid  and  following  P  57.  The  NaOH  reagent  should  also  be  tested 
for  sulfate  by  acidifying  it  with  HCl  and  adding  BaCl2,  since  the  presence  of 
much  sulfate  obscures  the  test  for  chromium  in  P.  55. 

6 


72  ANALYSIS  OF  THE  ALUMINUM-GROUP.  P.  64 

Procedure  54. — Confirmatory  Test  for  Aluminum. — Dissolve  the  pre- 
cipitate (P.  53),  or  a  small  portion  of  it  if  it  is  large,  in  5  cc.  HNOa. 
To  the  solution  add  from  a  dropper  half  as  many  drops  of  a  Co(NOa)j 
solution  containing  1%  Co  as  the  number  of  milligrams  of  aluminum 
estimated  to  be  present  in  the  solution.  Evaporate  the  solution 
almost  to  dryness  in  a  casserole,  add  a  drop  or  two  of  water,  and 
soak  up  the  solution  in  a  small  piece  of  filter  paper.  Make  a  small 
roll  of  the  paper,  wind  a  platinum  wire  around  it  in  the  form  of  a 
spiral,  and  heat  the  paper  in  a  flame  till  the  carbon  is  burnt  off.  (Blue 
residue,  presence  of  ALUMINUM.) 

Notes. — 1.  A  confirmatory  test  for  aluminum  should  always  be  tried  when 
the  NEUOH  precipitate  is  small;  for  the  precipitation  by  NH4OH  of  an  element 
whose  hydroxide  is  soluble  in  NaOH  is  not  very  characteristic  (lead,  antimony, 
tin,  and  silicon  showing  a  similar  behavior).  It  is  especially  necessary  to  guard 
against  mistaking  SiOsH^  for  A1(OH)3;  for  the  former  substance,  if  not  entirely 
removed  in  the  process  of  preparing  the  solution,  may  appear  at  this  point. 

2.  The  confirmatory  test  with  Co(NC>3)2  depends  upon  the  formation  of 
a  blue  substance,  whose  formula  is  not  definitely  known;  but  it  is  doubtless 
a  compound  of  the  two  oxides  CoO  and  Al2Oa,  and  is  probably  cobalt  alumi- 
nate,^Co(AlO2)2-  It  enables  0.5  mg.  Al  to  be  detected,  or  even  0.2  mg.  after 
a  little  practis.  No  other  element  givs  a  blue  color  to  the  ash.  It  is  essential 
to  have  the  aluminum  present  in  excess;  for  otherwise  the  blue  color  is  ob- 
scured by  the  black  oxide  of  cobalt.  Moreover,  when  sodium,  or  potassium 
salts  are  present,  the  ash  fuses  together  and  the  test  is  unsatisfactory.  For 
this  reason  the  sodium  salts  present  should  be  completely  washed  out  of  the 
NH40H  precipitate  before  dissolving  it  in  HNOj. 

Procedure  55. — Detection  of  Chromium. — To  the  NH4OH  solution  (P. 
53)  add  HAc,  1  cc.  at  a  time,  till  the  solution  reddens  blue  litmus  paper. 

If  the  solution  is  colorless,  treat  it  by  P.  57. 

If  it  is  at  all  yellow,  add  about  10  'cc.  BaCl2  solution,  and  heat  the 
mixture  to  boiling.  (Yellow  precipitate,  presence  of  CHROMIUM.) 
Filter,  and  wash  the  precipitate.  (Precipitate,  P.  56;  filtrate,  P.  57.) 

Notes. — 1.  The  presence  of  less  than  0.5  mg.  chromium  as  chromate  in  a 
volume  of  50  cc.  makes  the  solution  distinctly  yellow;  and  the  addition  of 
BaCl2  is  therefore  unnecessary  when  the  solution  is  perfectly  colorless.  It  is 
to  be  avoided,  since  BaSOi  may  be  precipitated  and  has  then  to  be  removed 
by  filtration.  In  doubtful  cases  the  color  of  the  solution  should  be  compared 
with  that  of  water.  The  color-test  is,  of  course,  not  delicate  by  artificial  light. 

2.  Since  some  sulfate  may  be  present,  the  formation  of  a  white  precipitate 
with  BaCl2  does  not  prove  the  presence  of  chromium.  Whether  the  precipitate 
is  pure  white  or  yellow  should  therefore  be  carefully  noted.  The  yellow  color 
of  a  small  BaCrO4  precipitate  is  most  apparent  when  the  precipitate  has  settled 
or  when  it  has  been  collected  on  the  filter.  If  there  be  sufficient  sulfate  present 
to  obscure  the  yellow  color  of  a  little  BaOC>4,  the  confirmatory  test  described  in 
the  next  procedure  will  enable  the  chromium  to  be  detected. 


P.  66  ANALYSIS  OF  THE  ALUMINUM-GROUP.  73 

Procedure  56. — Confirmatory  Test  for  Chromium. — Pour  repeatedly 
through  the  filter  (P.  55)  a  warm  mixture  of  3  cc.  HC1  and  10  cc.  sat- 
urated SO2  solution.  Evaporate  the  filtrate  in  a  casserole  just  to 
dryness,  taking  care  not  to  over-heat  the  residue ;  and  add  a  few  drops 
of  water.  (Green  coloration  of  the  solution,  presence  of  CHKOMIUM.) 

Note. — The  green  color  of  CrCla  is  intense  enough  to  enable  less  than  0.5 
mg.  of  chromium  in  1  cc.  of  solution  to  be  detected.  A  yellow  color  may  result 
from  the  action  of  the  acids  on  the  filter  paper;  but  this  color  usually  disappears 
on  evaporating  and  dissolving  the  residue  in  water. 

Procedure  57. — Detection  of  Zinc. — Warm  the  HAc  solution  (P.  55) 
to  50°  or  60°,  saturate  it  in  a  small  flask  with  H2S,  cork  the  flask, 
and  allow  it  to  stand  for  5  or  10  minutes  if  no  precipitate  separates 
at  once.  (White  flocculent  precipitate,  presence  of  ZINC.) 

Confirmatory  Test  for  Zinc. — Filter  the  H2S  precipitate  through  a 

double  filter  (made  by  folding  two  niters  together),  wash  it  with  a 

little  water,  and  pour  a  5-10  cc.  portion  of  HN03  repeatedly  through 

the  filter.     To  the  filtrate  add  from  a  dropper  one  drop  of  a  Co(N03)2 

solution  containing  1%  Co  and  as  many  more  drops  of  the  same 

solution  as  there  were  estimated  to  be  centigrams  of  zinc  in  the  H2S 

precipitate.     Evaporate  the  mixture  in  a  casserole  just  to  dryness  to 

expel  the  acid,  add  1  cc.  Na2COs  solution  and  0.5  cc.  more  for  each 

centigram  of  zinc  estimated  to  be  present.     Evaporate  to  dryness, 

ignite  gently  by  keeping  the  dish  moving  to  and  fro  in  a  small  flame 

until  the  purple  color  due  to  the  cobalt  disappears,  and  allow  the 

casserole  to  cool.     (Green  color,  presence  of  ZINC.)     [If  the  ignited 

mass  becomes  black  (owing  to  too  strong  heating),  add  a  few  drops 

HN03,  evaporate  just  to  dryness,  add  the  same  quantity  of  Na2COs 

solution  as  was  added  before,  evaporate  and  ignite  gently  as  before.] 

Notes. — 1.    ZnS  precipitates  more  rapidly,  and  in  a  somewhat  more  flocculent 

form,  from  a  warm  solution.    Very  small  quantities  of  zinc  (less  than  1  mg.) 

may  be  missed  unless  a  short  time  be  allowed  for  the  precipitate  to  coagulate; 

but  since  sulfur  may  then  separate,  the  appearance  of  a  white  turbidity  is  not 

sufficient  proof  of  the  presence  of  zinc.     The  precipitate  may  be  allowed  to 

settle  in  order  that  the  amount  of  zinc  present  may  be  better  estimated.    A 

double  filter  is  used,  since  the  ZnS  is  apt  to  pass  through  the  filter. 

2.  The  immediate  formation  of  a  white  flocculent  precipitate  with  H2S  in 
acetic  acid  solution  is  so  characteristic  as  to  be  a  sufficient  test  for  zinc.  Man- 
ganese is  the  only  other  element  of  this  group  that  forms  a  light-colored  sulfide; 
and  this,  owing  to  its  greater  solubility  in  water,  does  not  precipitate  from  an 
acetic  acid  solution.  The  confirmatory  test  described  in  the  last  paragraph  of 
the  procedure  is,  however,  useful  when  only  a  small  noncoagulating  precipitate 
which  may  be  sulfur  results,  or  when,  owing  to  the  presence  of  a  small  quantity 
of  other  elements,  the  precipitate  is  dark-colored.  In  this  test  the  use  of  an 
excess  of  Co(NOa)2  is  avoided,  since  otherwise  the  black  color  of  the  CoO  ob- 
scures the  green  color  of  the  cobalt  zincate  (CoZnO2). 


74 


ANALYSIS  OF  THE  IRON-GROUP. 


P.  61 


ANALYSIS    OF    THE     IRON-GKOUP. 


TABLE  IX. — ANALYSIS  OP  THE  IRON-GROUP. 


PRECIPITATE  PRODUCED  BY  SODIUM  HYDROXIDE  AND  PEROXIDE: 

A.  Phosphate  absent:  MnO(OH)2,  Fe(OH)3,  Co(OH)3,  Ni(OH)2, 

B.  Phosphate  present:  Also  BaCO3,  SrCO3,  CaCO3,  MgC03,  FePO4,  Ca3(PO4)2,  etc. 

Dissolve  in  HNOS  and  H202,  evaporate,  heat  with  HN03  and  KClOs  (P.  61). 


Precipitate: 
MnO-j. 
Add  HNOZ 
and  bismuth 
peroxide 
(P.  62). 

Solution:  Test  a  portion  for  phosphate  with  (NH^MoOi  (P.  68). 
A.  Phosphate  absent:  add  NH*OH  (P.  64). 
B.  Phosphate  present:  add  NH^Ac  and  FeCk,  dilute,  boil  (P.  65). 

Precipitate: 
A.  Fe(OH)3. 
B.  Basic  ferric 
acetate  and 
FeP04. 

Filtrate:  add  NH4OH,  pass  in  HyS  (P.  66). 

Precipitate:  ZnS,  CoS,  NiS. 
See  Table  X. 

Filtrate: 
A.  NHi  salts.    Reject. 
B.  Ba,  Ca,  Sr,  Mg. 
Treat  by  P.  81. 

Violet  color: 
HMn04. 

*  This  precipitate  may  contain  all  the  zinc  when  elements  of  the  iron-group  are  present  in  large  quantity. 

Procedure  61. — Precipitation  of  Manganese. — Transfer  the  Na2O2 
precipitate  (P.  52)  to  a  casserole,  together  with  the  filter  if  neces- 
sary, and  add  5-15  cc.  HNO3.  If  there  is  still  a  residue,  heat  the  mix- 
ture nearly  to  boiling  and  add  3%  H202  solution,  a  few  drops  at  a  time, 
-  -  stirring  after  each  addition,  till  the  residue  is  dissolved.  Filter  to 
remove  the  paper,  and  evaporate  the  filtrate  almost  to  dryness.  Add 
10  cc.  16-normal  HNO3  and  about  1  g.  of  powdered  KC1O3,  and  heat 
to  boiling.  (Black  precipitate,  presence  of  MANGANESE.) 

If  there  is  no  precipitate,  evaporate  the  solution  almost  to  dryness, 
add  25  cc.  water,  and  treat  a  part  of  the  solution  by  P.  63  and  the 
remainder  by  P.  64  or  P.  65. 

If  there  is  a  precipitate,  add  to  the  mixture  10  cc.  16-normal  HN03, 
pour  it  into  a  200  cc.  conical  flask,  heat  it  to  boiling,  and  add  grad- 
ually (0.5  g.  at  a  time)  3  g.  powdered  KC103.  Grasp  the  neck  of  the 
flask  with  a  strip  of  stiff  paper,  and  keep  the  flask  in  motion  over  a 
small  flame  so  that  the  mixture  boils  gently  for  one  or  two  minutes. 
Filter  with  the  aid  of  suction  through  an  asbestos  filter,  made  by  plac- 
ing in  a  funnel  enough  glass-wool  to  form  a  wad  1  cm.  high,  tamping 
it  down  gently  with  the  finger,  and  pouring  through  it  a  suspension  of 
asbestos  in  water,  enough  to  form  an  asbestos  layer  about  |  cm.  thick. 
Add  to  the  filtrate  1  g.  KC103.  If  more  precipitate  separates,  boil  the 


P.  61  ANALYSIS  OF  THE  IRON-GROUP.  75 

mixture  gently  for  a  minute  or  two,  and  then  filter  it  through  the  same 
asbestos  filter.  Wash  the  precipitate  two  or  three  times  with  16-nor- 
mal  HNOs  which  has  previously  been  freed  from  oxides  of  nitrogen 
by  warming  it  with  a  little  KClOs,  and  treat  the  precipitate  by  P.  62. 
Evaporate  the  filtrate  to  about  5  cc.,  but  not  further,  dilute  it  with 
25  cc.  water,  and  treat  a  part  of  it  by  P.  63  and  the  remainder  by 
P.  64  or  P.  65. 

Notes. — 1.  Pure  concentrated  HNOs  does  not  dissolve  hyd rated  MnOj; 
but  it  may  do  so  in  the  presence  of  filter-paper,  whereby  the  HNOs  is  reduced 
to  lower  oxides.  The  action  is  rapid  in  the  presence  of  E^CV,  for  the  MnOj 
is  thereby  quickly  reduced  to  Mn(NOs)2  with  evolution  of  Og. 

2.  By  HClOs  in  HNOs  solution  (but  not  by  HNOs  alone)  manganous  salts 
are  rapidly  oxidized  to  hydrated  MnC>2  with  formation  of  chlorin  dioxide  (C102), 
which  escapes  as  a  yellow  gas. 

3.  The  separation  of  manganese  in  this  way  from  the  other  metals  of  this 
group  is  entirely  satisfactory  with  the  exception  that  a  small  quantity  of  iron 
(up  to  1  mg.)  may  be  completely  carried  down  with  a  large  quantity  (500  mg.) 
of  manganese. 

Procedure  62. — Confirmatory  Test  for  Manganese. — Pour  through 
the  filter  containing  the  KClOs  precipitate  (P.  61)  5  cc.  hot  HNOs  to 
which  4  or  5  drops  of  3%  H202  solution  have  been  added.  Collect 
the  filtrate  in  a  test-tube;  cool  it;  add  to  it  solid  bismuth  dioxide, 
0.1  g.  at  a  time,  till  some  of  the  brown  solid  remains  undissolved;  and 
let  the  solid  settle.  (Purple  solution,  presence  of  MANGANESE.) 

Notes. — 1.  This  confirmatory  test  for  manganese  is  usually  superfluous, 
since  the  precipitation  of  MnO2  by  HClOs  is  highly  characteristic. 

2.  In  the  presence  of  HNOs,  MnO2  is  reduced  by  H2O2  with  evolution  of 
oxygen  and  formation  of  Mn(NOs)2. 

3.  Commercial  bismuth  dioxide,  also  often  called  sodium  bismuthate,  is 
a  mixture  of  bismuth  compounds  which  probably  owes  its  oxidizing  power  to 
the  presence  of  the  dioxide  Bi02.     When  it  is  not  available,  PbC>2  may  be 
substituted  for  it;  but  in  that  case  the  mixture  must  be  boiled  for  2  or  3  minutes. 

Procedure  63. — Test  for  Phosphate. — Pour  about  one-tenth  of  the 
HNOs  solution  (P.  61)  into  two  or  three  times  its  volume  of  ammonium 
molybdate  reagent,  and  heat  the  mixture  to  60-70°.  (Fine  yellow 
precipitate,  presence  of  PHOSPHATE.)  If  there  is  no  precipitate,  or 
only  a  very  small  one,  treat  the  remainder  of  the  HNOs  solution  by 
P.  64;  otherwise  by  P.  65. 

Notes. — 1.  Phosphate  is  tested  for  at  this  point  because  a  different  treatment 
is  necessary  when  it  is  present  in  significant  amount,  in  order  to  separate  from 
it  alkaline-earth  elements  and  to  provide  for  their  detection.  When  phosphate 
is  not  present,  iron  can  be  separated  from  nickel  and  cobalt  by  NH4<DH  (as  in 
P.  64);  but  when  considerable  phosphate  is  present,  alkaline-earth  elements 


76  ANALYSIS  OF  THE  IRON-GROUP.  P.  63 

may  also  be  present,  and  these  would  be  partly  or  wholly  precipitated  by 
NH4OH  as  phosphates.     (See  P.  51,  Note  6.) 

2.  In  order  that  the  phosphate  test  may  be  delicate  and  may  appear  imme- 
diately, a  large  proportion  of  the  molybdate  reagent  must  be  used  and  the 
solution  must  be  warmed.  The  precipitate  produced  by  ammonium  molybdate, 
(NH4)2MoO4,  is  ammonium  phospho-molybdate,  a  complex  salt  of  the  com- 
position (NH4)3PO4.12Mo03. 

Procedure  64. — Precipitation  of  Iron  in  Absence  of  Phosphate. — 
If  phosphate  is  absent,  make  the  HNOs  solution  (P.  61)  strongly 
alkaline  with  NH4OH*  using  an  excess  of  3-5  cc.  (Dark  red  precipi- 
tate, presence  of  IKON.)  Filter,  and  wash  the  precipitate  thoroly. 
Treat  the  nitrate  by  P.  66.  Dissolve  the  precipitate  in  HC1,  warming 
if  necessary;  and  to  the  solution  add  3-5  cc.  KCNS  solution.  (Red 
coloration,  presence  of  IRON.) 

Note. — This  test  may  be  made  in  the  presence  of  much  HC1,  but  not  in  the 
presence  of  much  HNOs;  for  HNOs  acts  on  KCNS  forming  a  red-colored  com- 
pound. This  test  for  iron  is  an  extremely  delicate  one;  and  if  only  a  faint 
color  is  obtained,  the  acids  used  in  the  process  must  be  tested  for  iron.  The 
red  color  is  due  to  the  formation  by  metathesis  of  ferric  thiocyanate,  Fe(CNS)s, 
a  slightly  ionized  substance. 

Procedure  65. — Detection  of  Iron  and  Removal  of  Phosphate  when 
Present. — If  phosphate  is  present,  test  one-tenth  of  the  HNOs  solu- 
tion (P.  61)  for  iron,  by  evaporating  it  just  to  dryness,  adding  1-2  cc. 
12-normal  HC1,  evaporating  again  just  to  dryness,  dissolving  the  resi- 
due in  2  cc.  HC1  and  5  cc.  water,  and  adding  5  cc.  KCNS  solution. 
(Red  coloration,  presence  of  IRON.)  To  the  remainder  of  the  solution 
add  NH4OH  until  the  precipitate  formed  by  the  last  drop  does  not  re- 
dissolve  on  shaking.  If,  owing  to  the  addition  of  too  much  NH4OH, 
the  solution  becomes  alkaline  or  a  large  precipitate  separates,  make 
it  distinctly  acid  with  acetic  acid.  Add  15  cc.  of  a  3-normal  solution 
of  NH4Ac,  and,  unless  the  mixture  is  already  of  a  brownish-red  color, 
add  Feds  solution,  drop  by  drop,  until  such  a  color  is  produced. 
Add  enough  water  to  make  the  volume  about  100  cc.,  boil  in  a  250  cc. 
flask  for  5  minutes,  adding  more  water  if  a  very  large  precipitate 
separates,  and  let  the  mixture  stand  for  a  minute  or  two.  Filter 
while  still  hot,  and  wash  with  hot  water.  Add  10  cc.  more  NH4Ac 
solution  to  the  nitrate,  boil  it  again,  and  collect  on  a  separate  filter 
any  further  precipitate.  Reject  the  precipitate.  Make  the  filtrate 
alkaline  with  NH4OH,  and  treat  it  by  P.  66. 

Notes. — 1.     With  regard  to  the  test  for  iron  with  KCNS  and  the  necessity 

of  removing  the  HNOs  by  the  evaporation  with  HC1,  see  P.  64,  Note. 

2.    This  method  of  separation  depends  on  the  facts  that,  upon  boiling  an 

acetic  acid  solution  containing  much  acetate,  ferric  iron  is  completely  precipi- 


P.  66  ANALYSIS  OF  THE  IRON-GROUP.  77 

tated  in  the  form  of  a  basic  acetate;  and  that  all  the  phosphate  present  com- 
bines with  the  iron  when  it  is  present  in  excess,  and  therefore  then  passes 
completely  into  the  precipitate,  leaving  the  bivalent  elements  in  solution. 
This  behavior  of  the  phosphate  is  due  to  the  fact  that  the  solubility  in  acids 
of  the  phosphates  of  the  trivalent  elements  is  much  smaller  than  that  of  the 
phosphates  of  the  bivalent  elements. 

3.  If  upon  adding  the  ammonium  acetate  the  solution  becomes  of  a  reddish 
color,  it  shows  that  iron  is  present  in  quantity  more  than  sufficient  to  combine 
with  the  phosphate;  for  a  cold  solution  containing  ferric  acetate  is  of  a  deep  red 
color.    If,  on  the  other  hand,  a  colorless  solution  results  (either  with  or  without 
a  precipitate),  it  shows  that  there  is  no  excess  of  iron,  and  FeCla  is  therefore 
added.     This  causes  the  precipitation  of  FeP(>4  as  a  yellowish  white  precipitate. 
Upon  boiling,  the  excess  of  iron  separates  completely  as  a  dark  red  gelatinous 
precipitate  of  basic  ferric  acetate,  leaving  the  supernatant  liquid  colorless, 
except  when  nickel  or  cobalt  is  present. 

4.  The  solution  is  diluted  to  at  least  100  cc.,  owing  to  the  large  volume  of 
the  precipitate;   and  it  is  heated  in  a  capacious  flask,  owing  to  its  tendency 
to  boil  over. 

Procedure  66. — Precipitation  of  Zinc,  Nickel,  and  Cobalt. — Into  the 
ammoniacal  solution  (P.  64  or  P.  65)  pass  H^S  gas  until  the  mixture 
after  shaking  blackens  PbAc2  paper  held  above  it.  (Black  precipi- 
tate, presence  of  NICKEL  or  COBALT.)  Filter,  wash  the  precipitate  with 
water  containing  a  very  little  (NH^S,  and  treat  it  by  P.  67.  Treat 
the  filtrate  by  P.  81-89  if  phosphate  or  much  chromium  was  found 
present  (in  P.  63  or  P.  55);  otherwise,  reject  it. 

Notes. — 1.  In  precipitating  NiS,  the  use  of  H2S  has  the  advantage  that  the 
nickel  is  all  thrown  down  at  once,  while  with  (NKi^S  some  of  it  usually  remains 
in  the  solution,  giving  it  a  dark  brown  color.  If  (NKi^S  be  used,  the  filtrate 
must  be  boiled  to  throw  down  the  unprecipitated  nickel,  as  described  in  P.  51. 
2.  The  filtrate  must  be  treated  by  the  procedures  for  detecting  alkaline- 
earth  elements  when  phosphate  has  been  found  to  be  present,  since  the  whole 
quantity  of  these  elements  present  in  the  substance  may  then  be  contained  in 
this  filtrate,  for  the  reasons  stated  in  Note  6,  P.  51,  and  Note  2,  P.  65.  The 
filtrate  should  also  be  so  treated  when  much  chromium  is  present,  since  all 
the  magnesium  present  may  have  been  carried  down  with  it  in  the  original 
precipitate. 


78  ANALYSIS  OF  THE  IRON-GROUP.  P.  67 

TABLE  X. — SEPARATION  OF  ZINC,  NICKEL,  AND  COBALT. 

HYDROGEN  SULPHIDE  PRECIPITATE  I  ZnS,  NiS,  CoS. 

Treat  with  cold  dilute  HCl  (P.  67}. 


Solution:  ZnCl2,  NiCl2*,  CoCl2* 

Add  NaOH  and  Na202  (P.  67). 


Filtrate:  Na2ZnO2. 
Add  HAc  and 


White  precipitate:  ZnS. 


Precipitate:  Ni(OH)2,  Co(OH)3, 
Add  HCl,  evaporate  (P.  68). 


Residue:  NiS,  CoS. 
Dissolve   in  HCl   and 

HNOs,   evaporate 

(P.  68). 


Residue :  NiCl2,  CoCl2.    Add  HCl  and  ether  (P.  68) . 


Yellow  residue:  NiCl2. 
Dissolve  in  water,  add  tartaric 
acid,  NaOH,  and  HzS  (P.  69). 


Brown    coloration:    presence 
of  nickel. 


Blue  solution:  CoCl2. 
Evaporate,  add  HAc  and 
KN02  (P.  70). 


Yellow  precipitate: 
K3Co(NO2)6. 


*  A  small  proportion  of  the  nickel  and  cobalt  present  always  dissolves  in  the  dilute  HCL 

Procedure  67. — Separation  of  Zinc  from  the  Nickel  and  Cobalt. — 
Transfer  the  H2S  precipitate  (P.  66)  with  the  filter  to  a  casserole, 
and  add  10-30  cc.  1-normal  HCl.  Stir  the  cold  mixture  frequently 
for  5  minutes,  and  filter.  Wash  the  residue  and  treat  it  by  P.  68. 

Boil  the  HCl  solution  until  the  H2S  is  completely  expelled,  add 
NaOH  solution  until  the  mixture  is  slightly  alkaline,  transfer  to  a 
casserole,  cool,  and  add  0.5-1  cc.  Na2C>2  powder,  a  small  portion  at 
a  time.    Boil  for  several  minutes  to  decompose  the  excess  of  Na202, 
cool  the  mixture,  and  filter.     Wash  the  precipitate  thoroly,  and  treat 
it  by  P.  68,  uniting  it  with  the  sulfide  residue  undissolved  by  the 
dilute  HCl.    Acidify  the  filtrate  with  HAc,  and  test  it  for  ZINC  by  P.  57. 
Notes. — 1.    This  treatment  with  dilute  HCl  serves  to  extract  almost  com- 
pletely the  zinc  which  may  be  present  in  this  precipitate  owing  to  its  having 
been  carried  down  in  the  Na202  precipitate,  as  described  in  P.  52,  Note  6. 
A  small  proportion  of  the  nickel  and  cobalt  present  (5-20%)  always  dissolves 
in  the  dilute  HCl,  and  the  subsequent  treatment  with  Na^O2  serves  to  separate 
these  elements  from  the  zinc.    This  separation  is  satisfactory  when,  as  in  this 
case,  the  nickel  and  cobalt  are  present  in  small  quantity;   for  then  only  an 
insignificant  amount  of  zinc  is  carried  down  with  them.    When,  therefore,  the 
H2S  precipitate  is  small,  it  may,  instead  of  being  treated  with  dilute  HCl,  be 
dissolved  at  once  in  aqua  regia  and  the  solution  treated  directly  as  described 
in  the  last  paragraph  of  this  Procedure. 

2.  This  Procedure  must  always  be  followed  in  order  to  determin  whether 
or  not  zinc  is  present  in  the  substance,  unless  a  satisfactory  test  for  it  has 
already  been  obtained  in  P.  57,  or  unless  the  original  Na-jC^  precipitate  (P.  52) 


P.  68  ANALYSIS  OF  THE  IRON-GROUP.  79 

was  small.    In  either  of  these  two  cases  this  Procedure  may  be  omitted  and 
the  H2S  precipitate  (P.  66)  treated  directly  by  P.  68. 

3.  The  fact  that  NiS  and  CoS  do  not  dissolve  readily  in  1-normal  HC1  seems 
inconsistent  with  the  non-precipitation  of  nickel  and  cobalt  by  B^S  with  the 
copper  and  tin  groups  from  a  solution  which  is  only  0.3  normal  in  acid.  This 
behavior  probably  arises  from  the  fact  that  these  sulfides  exist  in  at  least  two 
allotropic  forms  of  different  solubilities.  The  form  that  tends  to  be  produced 
by  direct  union  of  sulfide-ion  with  nickel-ion  or  cobalt-ion  is  soluble  in  dilute 
acid.  The  precipitate  produced  by  (NH^S  consists,  however,  only  in  part  of 
this  soluble  form:  it  contains  also  a  large  proportion  (increasing  with  the  time 
that  the  precipitate  has  stood)  of  a  form  that  is  nearly  insoluble  in  dilute  acids. 

Procedure  68. — Separation  of  Nickel  and  Cobalt. — Transfer  the 
sulfide  residue  undissolved  by  dilute  HC1  and  the  thoroly  washed  Na2O2 
precipitate  (P.  67)  with  the  filters  to  a  casserole,  add  5-15  cc.  HC1 
and  a  few  drops  HNOs,  warm  until  the  black  precipitates  are  dissolved, 
and  filter  off  the  paper.  Evaporate  the  solution  just  to  dryness 
moisten  the  residue  with  12-normal  HC1,  and  evaporate  again  just 
to  dryness.  To  the  dry  residue  add  0.5-2.0  cc.  12-normal  HC1; 
and,  if  it  does  not  all  dissolve,  warm  it  gently  with  constant  stirring. 
Cool,  add  10  cc.  ether  previously  saturated  with  HC1  gas,  and  stir  for 
1-3  minutes.  Filter  the  mixture  through  a  dry  filter,  and  wash  out 
the  filter  with  5-10  cc.  of  the  HCl-ether  reagent.  (Yellow  residue, 
presence  of  NICKEL;  blue  solution,  presence  of  COBALT.)  (Residue, 
P.  69;  solution,  P.  70.) 

Notes. — 1.  The  nickel  and  cobalt  must  be  present  as  chlorides  in  order  to  be 
separated  by  this  process.  The  two  evaporations  with  an  excess  of  HC1  serve 
to  destroy  any  nitrate  present. 

2.  Even  500  mg.  of  either  nickel  or  cobalt  dissolve  in  2  cc.  12-normal  HC1 
on  gentle  warming;  and  the  residue  should  be  so  dissolved  before  adding  the 
HCl-ether  reagent.    500  mg.  of  cobalt  will  then  remain  in  solution  on  the  addi- 
tion of  10  cc.  of  that  reagent. 

3.  Even  0.5  mg.  of  nickel  yields  a  distinct  yellow  residue;  and  less  than  0.5 
mg.  of  cobalt  imparts  a  pronounced  blue  coloration  to  the  ethereal  solution.     It 
is  therefore  unnecessary  to  try  the  confirmatory  test  for  cobalt  when  the  ethereal 
solution  is  colorless. 

4.  CoCb  in  aqueous  solution  is  pink,  like  Co(NC>3)2  or  CoSO4,  when  its 
concentration  is  small;  but  it  is  blue,  even  in  the  cold,  when  its  concentration 
is  very  large.     It  is  blue,  even  when  its  concentration  is  only  moderately 
large,  in  hot  solutions,  or  in  cold  solutions  which  contain  HC1  or  other  chlorides 
at  a  high  concentration.     It  is  also  blue  in  organic  solvents.    The  pink  color 
is  probably  due  to  the  cobalt-ion,  which  is  doubtless  hydrated  (existing  per- 
haps as  Co(H2O)4++  in  analogy  with  Co(NH3)4++),  and  to  the  hydrated  un- 
ionized cobalt  chloride  which  results  directly  from  the  union  of  this  ion  with 
chloride-ion.    The  blue  color  is  probably  due  to  anhydrous  CoCk,  or  to  com- 
plex anions,  like  Cods",  which  result  from  the  union  of  it  with  chloride-ion. 


80  ANALYSIS  OF  THE  IRON-GROUP.  P.  69 

Procedure  69. — Confirmatory  Test  for  Nickel. — Pour  through  the 
filter  containing  the  residue  undissolved  by  ether  (P.  68)  10  cc.  water, 
add  3-5  cc.  10%  tartaric  acid  solution,  neutralize  with  NaOH  solu- 
tion, and  add  2  cc.  in  excess.  Pass  in  H2S  gas  for  about  1  minute, 
filter  out  any  precipitate  that  may  form,  and  saturate  the  filtrate 
with  H2S.  Filter  again  if  there  is  a  precipitate.  (Brown  coloration, 
presence  of  NICKEL.) 

Notes. — 1.  When  an  alkaline  tartrate  solution  containing  a  very  little  nickel 
(even  0.1-0.2  mg.  in  20  cc.)  is  saturated  with  H2S,  a  clear  brown  solution  is 
obtained.  With  somewhat  larger  amounts  of  nickel  (10-20  mg.)  the  liquid  is 
opaque,  but  runs  through  a  filter  very  readily.  The  condition  of  the  nickel 
in  this  solution  is  not  known.  The  presence  of  the  tartrate  serves  merely  to 
prevent  the  precipitation  of  Ni(OH)2  by  the  NaOH  solution,  owing  to  the 
formation  of  a  complex  salt  containing  the  nickel  in  the  anion.  The  brown 
color  does  not  appear  until  the  alkaline  solution  is  nearly  saturated  with  H2S, 
so  that  care  must  be  taken  to  use  an  excess  of  H2S. 

2.  This  confirmatory  test  for  nickel  is  not  interfered  with  by  moderate 
amounts  of  other  elements  of  this  group,  such  as  cobalt  and  iron;  for  on  leading 
H2S  into  an  alkaline  tartrate  solution  containing  these  elements,  they  are 
completely  precipitated  as  sulfides  and  may  be  filtered  off,  yielding  a  filtrate 
which,  when  saturated  with  H2S,  remains  colorless  if  nickel  is  absent,  but  be- 
comes dark  brown  if  it  is  present  in  even  small  amount. 

Procedure  70. — Confirmatory  Test  for  Cobalt. — Heat  the  ethereal 
solution  (P.  68)  in  a  casserole  on  a  steam-bath,  or  by  floating  the 
casserole  in  a  vessel  of  nearly  boiling  water,  till  the  ether  is  expelled; 
then  evaporate  it  just  to  dryness  over  a  small  flame.  To  the  residue 
add  about  5  cc.  water,  and  a  few  drops  of  NaOH  solution.  To  the 
mixture  add  5  cc.  HAc  and  then  30  cc.  3-normal  KNO2  solution;  and 
allow  the  mixture  to  stand  at  least  ten  minutes  if  no  precipitate  forms 
sooner.  (Fine  yellow  precipitate,  presence  of  COBALT.) 

Notes. — 1.  The  precipitate  is  potassium  cobaltic  nitrite,  3KN02.Co(NO2)s, 
or  more  properly,  potassium  cobaltinitrite,  K3Co(NO2)e,  since  in  solution  it 
dissociates  into  K+  and  the  complex  anion  CotNO^e^.  In  the  formation 
of  this  substance  the  cobaltous  salt  is  oxidized  to  the  cobaltic  state  by  the 
nitrous  acid  displaced  from  its  salt  by  the  acetic  acid,  the  cobaltic  salt  combining 
as  fast  as  formed  with  the  potassium  nitrite. 

2.  The  precipitate  is  somewhat  soluble  in  water,  but  very  difficultly  soluble 
in  a  concentrated  KNO2  solution,  owing  to  the  common-ion  effect  of  the  potas- 
sium-ion.   The  formation  of  the  KsCo(NO2)6  precipitate  takes  place  slowly;  but, 
even  when  only  0.1-0.2  mg.  of  cobalt  is  present,  a  distinct  precipitate  results 
within  10  minutes. 

3.  Nickelous  salts  are  not  oxidized  by  nitrous  acid,  and  they  are  not  precip- 
itated by  KNO2,  unless  a  very  large  quantity  of  nickel  is  present,  in  which 
case   a  dark-yellow  or  dark-red   precipitate  of  potassium   nickelous   nitrite, 
K4Ni(NO2)e,  may  separate. 


P.  81  ANALYSIS  OF  ALKALINE-EARTH  GROUP.  81 

PRECIPITATION    AND    ANALYSIS    OF    THE    ALKALINE-EARTH    GROUP. 


TABLE  XI.  —  ANALYSIS  OF  THE  ALKALINE-EARTH  GROUP. 


AMMONIUM  CARBONATE  PRECIPITATE:  BaCO3,  SrCO3,  CaCO3, 
Dissolve  in  HAc,  add  NH+Ac  and  KzCrO*  (P.  82}. 


Precipitate: 
BaCrO4. 
Dissolve  in  HCl, 
evaporate  (P.  88). 

Filtrate.    Add  NHtOH  and  alcohol  (P.  84). 

Precipitate: 
SrCr04. 
Treat  with 
(NH^COZ  (P.  85). 

Filtrate:    Ca  and  Mg  salts. 
Add  (ATtf4)2C204  (P.  86). 

Test  in 
flame 

Add  HAc, 
NHtAc, 
and 
KzCrO*. 

Precipitate: 
CaQA. 

Dissolve  in  dilute 
H2SOt,  add  alcohol 
(P.  87). 

Filtrate. 
Add  NH&H 
and  Na2HPOi 
(P.  88). 

Residue:  SrCO3. 
Dissolve  in  HAc, 
add  CaSO*. 

Green 
color  : 
Ba. 

Precipi- 
tate: 
BaCrO4. 

Precipitate: 
MgNH4PO4. 

Precipitate  : 
SrSO4. 

Precipitate: 
CaS04. 

Procedure  81. — Precipitation  of  the  Alkaline-Earth  Group. — 
Evaporate  the  filtrate  from  the  NH4OH  and  (NH4)2S  precipitate 
(P.  51)  to  a  volume  of  about  10  cc.,  and  filter  off  the  sulfur. 

To  the  cold  solution  add  15  cc.  (NH4)2C03  reagent  and  15  cc. 
95%  alcohol;  and,  if  a  large  precipitate  results,  add  15  cc.  more  of 
each  of  these  liquids.  Shake  the  mixture  continuously  for  10  min- 
utes; or,  let  it  stand,  with  occasional  shaking,  for  at  least  half  an  hour. 
(Precipitate,  presence  of  ALKALINE-EARTH  ELEMENTS.)  Filter,  and 
wash  the  precipitate  with  a  little  (NH4)2CO3  reagent,  using  suction 
if  the  precipitate  is  large.  (Precipitate,  P.  82;  filtrate,  P.  91.) 

Notes. — 1.  The  nitrate  from  the  (NH4)2S  precipitate  is  evaporated  in  order 
that  the  elements  of  the  alkaline-earth  group  may  be  precipitated  more  quickly 
and  more  completely.  The  evaporation  also  serves  to  destroy  (NHU)2S  and  to 
coagulate  any  sulfur  that  may  separate.  The  volume  to  which  the  (NH^COs 
reagent  is  added  should  not  exceed  10  cc. 

2.  The  (NH4)2CO3  reagent  is  prepared  by  dissolving  250  g.  freshly  powdered 
ammonium  carbonate  in  1  liter  6-normal  NHiOH. 

3.  If  the  ammonium  carbonate  and  hydroxide  were  added  in  only  small 
excess,  the  precipitation  of  CaC03,  SrCO3,  and  BaCOs  would  not  be  complete, 
and  additional  tests  for  small  quantities  of  these  elements  would  have  to 
be  made  in  the  nitrate.    But,  by  the   use   of   a   concentrated   solution  of 
(NH4)2CO3  containing  a  large  excess   of   NBUOH   (so   as  to   diminish   the 
hydrolysis    of   the    carbonate   into    (NH^+HCOr    and    NHiOH),    the   pre- 


82  ANALYSIS  OF  ALKALINE-EARTH  GROUP.  P.  81 

cipitation  may  be  made  practically  complete,  owing  to  the  greatly  increased 
concentration  of  carbonate-ion  (COa"). 

4.  When  the  concentrations    of  (NH^COs  and   NH4OH  are  sufficiently 
great,  magnesium  is  in  the  cold  also  completely  precipitated.     The  precipitate, 
which  is  in  this  case  a  double  carbonate,  MgCOs^NH^COa^EkO,  is,  however, 
fairly  soluble  in  cold  water  and  readily  soluble  in  hot  water. 

5.  From  a  cold  aqueous  solution  the  precipitation  of  these  elements  takes 
place  slowly,  especially  in  the  case  of  magnesium  and  calcium;  but  it  is  greatly 
accelerated  by  the  addition  of  alcohol  and  by  shaking.    Under  the  conditions 
recommended  in  the  procedure  0.5  mg.  of  any  of  the  four  elements  is  easily 
detected. 

Procedure  82.— Precipitation  of  Barium— Dissolve  the  (NH4)2COS 
precipitate  (P.  81)  by  pouring  repeatedly  through  the  filter  a  5-15 
cc.  portion  of  hot  HAc,  and  evaporate  the  solution  just  to  dryness, 
taking  care  not  to  ignite  the  residue. 

Add  to  the  residue  3  cc.  HAc,  20  cc.  3-normal  NH^c,  and  15  cc. 
water;  and  heat  the  solution  to  boiling  in  a  flask.  Measure  out  8 
cc.  3-normal  K2Cr04  solution,  and  add  it  a  few  drops  at  a  time,  heat- 
ing and  shaking  after  each  addition.  Finally,  heat  the  mixture  at 
90-100°  for  1  or  2  minutes,  shaking  at  the  same  time.  Filter,  even 
tho  the  solution  appear  clear;  remove  the  filtrate,  and  wash  the 
precipitate  thoroly  with  cold  water.  (Pale  yellow  precipitate, 
presence  of  BARIUM.)  (Precipitate,  P.  83;  filtrate,  P.  84.) 

Notes. — 1.  The  solubility  in  water  of  the  chromates  of  the  alkaline-earth 
elements  increases  rapidly  in  the  order,  Ba,  Sr,  Ca,  Mg,  as  shown  in  the 
Table  on  page  124.  The  difference  in  solubility  of  BaCrO4  and  SrCrO4  is  so 
great  that  under  the  conditions  of  the  procedure  0.5  mg.  Ba  can  be  detected, 
while  even  400  mg.  Sr  giv  no  precipitate.  The  amount  of  K2CrO4  added  is 
sufficient  to  precipitate  completely  more  than  500  mg.  of  barium. 

2.  Acetic  acid  is  added  to  increase  the  solubility  of  SrCrC^.    By  its  action 
the  concentration  of  the  chromate-ion  is  decreased,  owing  to  its  conversion 
partly  into  hydrochromate-ion  and  partly  into  bichromate-ion,  according  to  the 
reactions: 

CrOr  +  H+  =  HCrO4-;  and  2HCrO4-  =  H2O  +  Cr2O7". 
It  is  evident  that  the  ratio  of  the  CrO4~  to  the  HCrO4~  concentration  must 
decrease  as  the  H+  concentration  increases.  For  this  reason  the  presence  of  an 
excess  of  a  largely  ionized  acid  (such  as  HC1  or  HNOa)  would  prevent  the  pre- 
cipitation of  BaCrO4;  but  since  acetic  acid  is  a  slightly  ionized  acid,  and  since 
a  large  amount  of  acetate  is  present,  the  addition  of  a  considerable  excess  of 
acetic  acid  has  but  little  effect. 

3.  The  K2Cr04  is  added  slowly  to  the  hot  solution  and  the  mixture  is  shaken 
and  heated  in  the  neighborhood  of  100°  before  filtering,  since  otherwise  the 
precipitate  is  liable  to  pass  through  the  filter.    By  this  method  of  precipitation 
almost  all  the  barium  is  precipitated  before  an  excess  of  K2CrO4  is  added. 
This  is  of  importance,  since,  when  much  barium  is  present,  as  much  as  3  mg. 


P.  83  ANALYSIS  OF  ALKALINE-EARTH  GROUP.  83 

Sr  may  be  carried  down  completely  if  the  K2(M)4  reagent  be  added  quickly. 
If  for  any  reason  the  filtrate  is  turbid  after  two  or  three  filtrations,  the  pre- 
cipitate may  be  coagulated  by  boiling  gently  for  1  or  2  minutes.  Vigorous 
or  long- continued  boiling  is  to  be  avoided,  since,  owing  to  loss  of  acetic  acid, 
SrCrCU  may  then  separate  if  much  strontium  is  present.  When  less  than  1 
mg.  Ba  is  present  it  is  very  difficult  to  distinguish  the  faint  turbidity  in  the 
colored  solution;  but  the  pale  yellow  precipitate  can  be  seen  after  filtering  and 
washing  the  K2CrO4  out  of  the  filter.  The  precipitate  must  be  washed  thoroly, 
in  order  to  remove  strontium  as  completely  as  possible,  which  otherwise  would 
obscure  the  flame  coloration  in  the  confirmatory  test  for  barium. 

Procedure  83. — Confirmatory  Test  for  Barium. — Pour  repeatedly 
through  the  filter  containing  the  K2Cr04  precipitate  (P.  82)  a  5-10  cc. 
portion  of  hot  HC1,  and  evaporate  the  solution  just  to  dryness  in  a 
small  casserole.  Heat  a  platinum  wire  having  a  small  loop  at  one 
end  in  a  gas  flame  till  it  no  longer  colors  the  flame,  dip  it  in  HC1, 
touch  it  to  the  residue  in  the  casserole,  and  introduce  it  again  into  the 
flame.  (Green  color,  presence  of  BARIUM.)  If  this  flame  test  is  in- 
conclusiv,  treat  the  residue  by  the  second  paragraph  of  P.  82.  (Yel- 
low precipitate,  presence  of  BARIUM.) 

Notes. — 1.  When  the  amount  of  barium  present  is  very  small,  only  a  mo- 
mentary green  color  is  seen  as  the  yellow  sodium  color  which  first  appears 
fades  away.  The  only  other  elements  that  color  the  flame  green  are  copper, 
boron,  and  thallium.  Strontium  givs  a  crimson  and  calcium  an  orange-red  color. 

2.  If  a  small  quantity  of  SrCrC>4  was  precipitated  in  P.  82,  owing  to  the  pres- 
ence of  a  very  large  amount  of  strontium  in  the  substance  or  owing  to  failure  to 
follow  the  directions,  it  will  not  again  precipitate  in  this  second  treatment; 
for  the  quantity  of  strontium  now  present  in  the  solution  is  much  less  than 
before.  A  yellow  precipitate  obtained  in  this  second  treatment  is  therefore 
conclusiv  evidence  of  the  presence  of  barium. 

Procedure  84. — Precipitation  of  Strontium. — To  the  filtrate  (P.  82), 
after  cooling  it,  add  NEUOH  slowly  until  the  color  of  the  solution 
changes  from  orange  to  yellow,  and  then  10  cc.  more.  Dilute  the 
solution  to  65  cc.,  and  add  slowly,  with  constant  shaking,  50  cc.  95% 
alcohol,  and  cool  the  solution.  If  a  large  precipitate  forms  or  if  a  large 
precipitate  of  BaCr04  separated  in  P.  82,  add  4  cc.  3-normal  K2Cr04 
solution  and  10  cc.  95%  alcohol,  and  shake.  (Light  yellow  precipitate, 
presence  of  STRONTIUM.)  Filter  after  5-10  minutes  with  the  aid  of 
suction;  suck  the  precipitate  as  dry  as  possible,  but  do  not  wash  it. 
(Precipitate,  P.  85;  filtrate,  P.  86.) 

Notes. — 1.  Under  these  conditions  1  mg.  of  strontium  yields  a  precipitate, 
while  even  400-500  mg.  of  calcium  or  magnesium  do  not  do  so.  A 
moderate  change  in  the  conditions  will  not  affect  this  result;  but  if  the  con- 
centration of  alcohol  or  K2CrO4  is  much  less  than  is  recommended,  the  pre- 
cipitation of  strontium  may  be  incomplete;  while  the  addition  of  larger  amounts 


84  ANALYSIS  OF  ALKALINE-EARTH  GROUP.  P.  84 

of  alcohol  and  K2CrO4  may  cause  the  precipitation  of  chromate  of  calcium  or 
magnesium  if  much  of  these  elements  is  present,  or  of  K2CrO4  itself,  since  the 
latter  is  not  very  soluble  in  alcohol.  The  confirmatory  test  should  therefore 
always  be  tried. 

2.  When  a  large  quantity  of  strontium  or  barium  is  present,  the  K2OO4 
added  in  P.  82  may  have  been  so  far  removed  from  the  solution  that  the  stron- 
tium is  not  at  first  completely  precipitated.    In  such  cases,  in  order  to  ensure 
its  complete  precipitation,  more  K2CrO4  is  added  in  this  Procedure. 

3.  The  precipitate  is  not  washed,  since  SrCrO4  is  fairly  soluble  in  water. 

Procedure   85. — Confirmatory  Test  for  Strontium. — If  the  K2Cr04 
precipitate  (P.  84)  is  small,  pour  three  or  four  times  through  the  filter 
containing  it  a  hot  mixture  consisting  of  10  cc.  (NH4)2C03  reagent 
and  5  cc.  3-normal  K2Cr04  solution.     If  the  K2Cr04  precipitate  is 
large,  transfer  it  to  a  casserole,  heat  it  for  4-5  minutes  with  15  cc.  of 
the  same  mixture,  and  filter.    Wash  the  residue  thoroly  with  cold 
water.    Pour  repeatedly  through  the  filter  containing  it  a  cold  5-10 
cc.  portion  of  1-normal  HAc.    Evaporate  the  solution  just  to  dryness, 
add  2-3  cc.  water,  transfer  the  solution  to  a  test-tube  (pouring  it 
through  a  very  small  filter,  if  it  is  not  perfectly  clear),  add  3  cc.  satu- 
rated CaS04.2H20  solution,  and  heat  in  a  vessel  of  boiling  water  for  at 
least  ten  minutes.     (Fine  white  precipitate,  presence  of  STRONTIUM.) 
Note.— The  mixture  of  (NH4)2CO3  and  K2CrO4  converts  SrCrO4  into  the 
less  soluble  SrCO^;  but  it  has  no  action  on  any  BaCrC>4  that  may  be  present, 
and  only  very  little  of  any  such  BaCrO4  dissolves  in  the  subsequent  treatment 
with  HAc.     Any  CaCrO4  present  is  converted  by  the  mixture  into  CaCOs, 
and  this  dissolves  in  the  HAc;   but  a  solution  of  a  calcium  salt  givs  no  pre- 
cipitate with  saturated  CaSO4  solution  (tho  it  might,  of  course,  do  so  with  a 
concentrated  solution  of  another  sulfate,  like  K2SO4). 

Procedure  86. — Precipitation  of  Calcium. — To  the  filtrate  from  the 
rOi  precipitate  (P.  84)  add  60  cc.  water,  heat  to  it  boiling,  add 
to  it  slowly  60  cc.  boiling-hot  (NH4)2C204  solution,  and  keep  it  nearly 
boiling  for  5  minutes.  (White  precipitate,  presence  of  CALCIUM.) 
Filter  the  mixture  while  it  is  still  hot;  and  wash  the  precipitate 
once  or  twice  with  water.  (Precipitate,  P.  87;  filtrate,  P.  88.) 

Note. — The  solution  is  heated  to  boiling,  and  the  ammonium  oxalate  solution 
is  added  slowly,  in  order  to  cause  the  CaC2O4  to  precipitate  in  the  form  of 
coarser  particles  which  can  be  more  readily  filtered.  Moreover,  the  solution 
must  be  heated  and  must  be  kept  hot,  in  order  to  prevent  the  precipitation 
of  some  MgC2C>4,  which  would  precipitate  when  much  magnesium  is  present 
if  the  mixture  were  allowed  to  cool. 

Procedure  87. — Confirmatory  Test  for  Calcium. — Pour  repeatedly 
through  the  filter  containing  the  (NH4)2C204  precipitate  (P.  86)  a  cold 
5  cc.  portion  of  H2S04;  add  10-15  cc.  95%  alcohol,  and  let  the  mixture 
stand  for  several  minutes.  (White  precipitate,  presence  of  CALCIUM.) 


P.  87  ANALYSIS  OF  ALKALINE-EARTH  GROUP.  85 

Notes. — 1.  CaC2O4.H2O  is  very  difficultly  soluble  in  water,  but  dissolves  in 
dilute  solutions  of  largely  ionized  acids,  owing  to  the  formation  by  metathesis  of 
unionized  HC2O4~.  CaSO4  is  somewhat  soluble  in  dilute  H2SO4,  but  is  com- 
pletely thrown  out  as  a  flocculent  precipitate  by  the  addition  of  the  alcohol. 

2.  One  milligram  of  calcium  produces  a  turbidity  at  once,  0.5  mg.  in  1-3 
minutes,  and  0.2  mg.  within  10  minutes.  Even  a  large  amount  of  magnesium 
does  not  interfere  with  the  test.  If  strontium  were  present,  a  small  amount  of  it 
would  dissolve  in  the  H2SO4,  but  only  enough  to  giv  a  slight  turbidity  on  the 
addition  of  alcohol,  corresponding  to  that  given  by  0.2-0.3  mg.  of  calcium  after 
standing  a  few  minutes.  Therefore  anything  more  than  a  slight  turbidity  ia 
a  conclusiv  proof  of  the  presence  of  calcium. 

Procedure  88. — Detection  of  Magnesium. — To  the  filtrate  from  the 
(NH4)2C204  precipitate  (P.  86)  add  10  cc.  15-normal  NH4OH  and 
25  cc.  Na2HP04  solution;  cool,  and  shake  the  mixture;  if  no  precipi- 
tate forms,  let  the  mixture  stand  for  at  least  half  an  hour,  shaking  it 
frequently.  (White  precipitate,  presence  of  MAGNESIUM.)  Filter  out 
the  precipitate,  wash  it  once  with  alcohol,  and  treat  it  by  P.  89. 

Note. — This  test  for  magnesium  depends  upon  the  precipitation  of  magne- 
sium ammonium  phosphate,  Mg(NEL|)PO4.  This  salt  is  fairly  soluble  even 
in  cold  water,  owing  chiefly  to  hydrolysis  into  NH4OH  and  Mg++HPO4~; 
and  the  test  is  therefore  made  in  a  strongly  ammoniacal  solution.  Since  the 
solubility  increases  rapidly  with  the  temperature,  the  solution  is  cooled  to  the 
room  temperature,  or  below.  In  an  aqueous  solution  this  substance  shows 
a  great  tendency  to  form  a  supersaturated  solution,  and  it  is  therefore  usually 
directed  to  make  the  test  in  as  small  a  volume  as  possible.  In  the  presence  of 
alcohol,  however,  precipitation  takes  place  rapidly,  and  even  0.5  mg.  of  mag- 
nesium produces  a  distinct  turbidity  in  a  solution  containing  190  cc.  water  and 
50  cc.  alcohol  within  half  an  hour.  A  small  precipitate  of  this  kind  settles  out  on 
further  standing  and  may  then  be  detected  by  rotating  the  solution  so  as  to 
cause  the  precipitate  to  collect  in  the  center. 

Procedure  89. — Confirmatory  Test  for  Magnesium. — Pour  repeatedly 
through  the  filter  containing  the  Na2HP04  precipitate  (P.  88)  a  5  cc. 
portion  of  2-normal  H2S04;  add  to  the  solution  20  cc.  95%  alcohol, 
and  shake  it  continuously  for  two  or  three  minutes.  Filter,  if  there 
is  a  precipitate;  add  to  the  filtrate  10  cc.  water,  20  cc.  NH4OH, 
and  5  cc.  Na2HPO4  solution;  and  let  the  mixture  stand  at  least  half  an 
hour.  (White  crystalline  precipitate,  presence  of  MAGNESIUM.) 

Notes. — 1.  This  confirmatory  test  should  be  tried  whenever  Na2HPC>4  pro- 
duces a  small  precipitate  that  is  not  distinctly  crystalline.  For,  if  even  a 
small  quantity  of  strontium  or  calcium  failed  to  be  precipitated  in  P.  84  or  86, 
a  flocculent  precipitate  of  Sra(P04)2  or  Ca3(PO4)2  would  come  down  on  the 
addition  of  Na^HPCU. 

2.  The  addition  of  H2SO4  and  alcohol  precipitates  strontium  and  calcium  so 
completely  that  any  precipitate  (more  than  a  very  slight  turbidity)  produced 
on  adding  Na2HPO4  to  the  HgSOi  solution  can  not  be  due  to  these  elements. 


86 


ANALYSIS  OF  THE  ALKALI-GROUP. 
ANALYSIS   OF   THE  ALKALI-GROUP. 


P.  91 


TABLE  XII. — ANALYSIS  OP  THE  ALKALI-GROUP. 


FILTRATE    FROM    AMMONIUM    CARBONATE    PRECIPITATE :       NIL;,  K,  Na  Salts. 

Evaporate  and  ignite  the  residue  (P.  91). 


Vapor: 
NH4  salts. 

Residue:  KC1,  NaCl.    Add  HCIO*,  evaporate,  add  alcohol  (P.  92). 

Residue:  KClOi. 
Dissolve  in  hot  water, 
addNazCo(NOz)s(P.9S). 

Solution:  NaClO4. 
Saturate  with  HCl  gas  (P.  94). 

Precipitate:  NaCl. 
Dissolve  in  water,  add 
K2H2Sbz07  (P.  95). 

Yellow  precipitate: 
K2NaCo(N02)6. 

Crystalline  precipitate: 
Na2H2Sb2O7. 

Procedure  91. — Removal  of  Sulfate  and  of  Ammonium  Salts. — 
Evaporate  the  filtrate  from  the  (NH4)2CO3  precipitate  (P.  81)  to  a 
volume  of  10-15  cc.  till  it  no  longer  smells  of  ammonia. 

To  10  drops  of  the  filtrate  add  4-5  drops  HCl  and  1-2  cc.  BaCl2 
solution.  (White  precipitate,  presence  of  SULFATE.) 

If  BaCl2  produces  no  precipitate,  treat  the  rest  of  the  filtrate  by 
the  last  paragraph  of  this  Procedure. 

If  BaCl2  produces  a  precipitate,  add  to  the  rest  of  the  filtrate  5-10 
cc.  BaCl2  solution,  heat  the  mixture  to  boiling,  and  filter  out  the  pre- 
cipitate. To  the  filtrate  add  5-15  cc.  (NH4)2CO3  reagent,  heat  the 
mixture  to  boiling,  filter  out  the  precipitate,  and  and  treat  the  fil- 
trate by  the  last  paragraph  of  this  Procedure. 

Evaporate  the  filtrate  to  dryness  in  a  small  casserole,  and  ignite 
the  residue,  at  first  gently,  then  just  below  redness,  until  no  more 
white  fumes  come  off,  keeping  the  dish  in  motion  over  a  flame  and 
taking  care  to  heat  the  sides  as  well  as  the  bottom  of  the  dish.  Cool 
completely,  add  5  cc.  water,  filter  the  mixture  through  a  5-cm.  filter, 
evaporate  the  filtrate  to  dryness  in  a  small  casserole,  and  heat  the 
Pish  at  a  temperature  just  below  redness.  (White  residue,  presence 
of  POTASSIUM  or  SODIUM.)  Treat  the  residue  by  P.  92. 

Notes. — 1.  If  sulfate  is  present,  either  because  it  was  a  constituent  of  the 
substance  or  because  H2SO4  was  used  in  the  preparation  of  the  solution  (in  P. 
5  or  P.  8),  it  must  be  removed  before  attempting  to  separate  potassium  and 
sodium  by  P.  92.  The  process  of  removing  it  involves  its  precipitation  as  BaSO4 
by  the  addition  of  BaCl2,  the  removal  of  the  excess  of  barium  by  the  subsequent 


P.  91  ANALYSIS  OF  THE  ALKALI-GROUP.  87 

addition  of  (NH^COs,  and  finally  the  volatilization  of  the  ammonium  salts  by 
ignition.  If  phosphate  is  present,  it  is  also  removed  by  the  addition  of  BaCk  to 
the  neutral  solution;  but,  as  it  is  not  necessary  to  remove  it,  the  preliminary 
test  with  BaCl2  is  made  in  HC1  solution,  in  which  case  sulfate  alone  produces 
a  precipitate. 

2.  Great  care  must  be  taken  to  volatilize  the  ammonium  salts  completely, 
since  even  1  mg.  of  ammonium  would  giv  a  precipitate  in  the  subsequent  test 
for  potassium  in  P.  92.     To  ensure  their  removal  the  residue  is  ignited  twice. 
The  dish  must  not,  however,  be  allowed  to  become  red-hot  during  the  ignition, 
since  at  that  temperature  potassium  and  sodium  chlorides  are  somewhat  volatil. 

3.  The  formation  of  a  precipitate  with  BaClz  does  not  necessarily  show 
that  sulfate  is  present  in  the  original  substance,  even  when  H2SO4  has  not 
been  used  in  preparing  the  solution;  for  it  may  have  been  produced  by  the 
oxidation  of  sulfide,  when  the  substance  contains  it,  in  preparing  the  solution 
in  P.  3  and  4;  or  it  may  even  be  formed  in  small  quantity  by  the  oxidation  of 
the  H2S  and  (NH^S  used  as  reagents. 

4.  A  brown  or  black  residue  of  organic  matter,  coming  from  impurity  in  the 
ammonium  salts  added  in  the  course  of  analysis  and  from  the  alcohol  and  filter 
paper,  may  remain  upon  treating  the  ignited  residue  with  water.     There  may 
also  be  a  white  residue  of  silica,  coming  from  the  action  of  the  reagents  on  the 
glass  and  porcelain  vessels  throughout  the  course  of  the  analysis. 

5.  Since  in  the  dry  form  even  a  residue  that  seems  very  small  may  correspond 
to  an  appreciable  quantity  of  potassium  or  sodium,  the  subsequent  tests  for  these 
elements  should  be  made  if  there  is  any  residue  whatever  after  the  final  ignition. 

Procedure  92. — Separation  of  Potassium  and  Sodium. — To  the  ig- 
nited residue  (P.  91)  add  5-15  cc.  2-normal  HC1O4;  and  evaporate,  by 
keeping  the  dish  in  motion  over  a  small  flame,  till  thick  white  fumes  of 
HC1O4  come  off  copiously  (which  occurs  when  the  liquid  is  reduced 
to  about  one-fourth  of  its  original  volume).  Cool  completely,  add 
10-20  cc.  95%  alcohol,  and  stir  the  mixture  for  2-5  minutes  if  there 
is  much  residue.  If  there  is  still  a  res 'due,  add  3  cc.  2-normal 
HC104.  (White  residue,  presence  of  POTASSIUM.)  Filter  through  a 
dry  filter-paper,  and  wash  the  residue  with  95%  alcohol.  (Precipi- 
tate, P.  93;  filtrate,  P.  94.) 

Notes. — 1.  Enough  HC1O4  must  be  added  to  convert  the  potassium  and 
sodium  chlorides  completely  into  perchlorates,  and  the  evaporation  must  be  con- 
tinued till  all  the  HC1  is  expelled;  for  otherwise  NaCl,  being  insoluble  in  alcohol, 
may  be  left  as  a  residue  with  the  KClOi.  An  unnecessary  excess  of  HC1O4  is, 
however,  to  be  avoided,  since  it  makes  the  subsequent  test  for  sodium  somewhat 
less  delicate.  The  quantity  added  is  therefore  varied  (from  5  to  15  cc.)  in  ac- 
cordance with  the  size  of  the  ignited  residue  obtained  in  P.  91. 

2.  The  evaporation  is  continued  till  the  HC1O4  fumes  strongly,  also  for  the 
purpose  of  removing  most  of  the  water;  for  this  test  for  potassium  and  the 
subsequent  test  for  sodium  (in  P.  94)  are  more  delicate,  the  less  the  quantity  of 
water  present.     When  the  directions  given  in  the  Procedure  are  followed,  1-1 J 
mg.  of  potassium  produces  a  distinct  precipitate. 

3.  If  sulfate  were  present  in  the  (NH4)2CO8  filtrate  and  it  were  not  removed 


88  ANALYSIS  OF  THE  ALKALI-GROUP.  P.  92 

in  P.  91  by  the  addition  of  BaCk,  this  separation  of  potassium  and  sodium  would 
be  unsatisfactory;  for,  when  sodium  is  present,  Na2SC>4  would  remain  in  the 
residue  undissolved  by  the  alcohol.  This  arises  from  the  fact  that  H2SO4  is  less 
volatil  than  HC1O4  and  is  therefore  not  expelled  by  it  in  the  evaporation,  and 
from  the  fact  that  Na2SC>4  is  only  slightly  soluble  in  alcohol  even  in  the  presence 
of  HC1O4.  The  presence  of  phosphate  or  borate  does  not,  however,  interfere 
with  the  separation;  for,  tho  phosphoric  and  boric  acids  are  not  volatilized 
in  the  evaporation  with  HC1O4  and  tho  their  sodium  salts  are  very  slightly 
soluble  in  pure  alcohol,  yet  these  salts  are  metathesized  by  the  excess  of  per- 
chloric acid,  since  this  acid  is  much  more  largely  ionized  than  phosphoric  or 
boric  acid.  The  sodium  therefore  remains  dissolved  in  the  alcoholic  solution. 

4.  A  small  quantity  of  sulfate  is  usually  produced  when  the  (NH4)2COj 
filtrate  is  evaporated  to  dryness  and  the  residue  is  ignited,  owing  to  the  action 
of  the  nitrate  present  on  sulfur-compounds  coming  from  the  decomposition  of 
the  (NH^S  reagent.  To  redissolve  any  precipitate  of  NazSC^  arising  from 
this  small  quantity  of  sulfate,  3  cc.  more  of  2-normal  HCICX  are  added  at  the 
end  of  the  Procedure  when  there  is  a  residue  undissolved  by  the  alcohol. 

Procedure  93. — Confirmatory  Test  for  Potassium. — Pour  repeatedly 
through  the  filter  containing  the  HC104  precipitate  (P.  92)  a  5-10  cc. 
portion  of  boiling  water,  add  to  the  solution  3-5  cc.  NaNOa  solution 
and  2-3  cc.  HAc,  and  boil  the  mixture  gently  in  a  casserole  for  5-6  min- 
utes.    Cool  the  mixture,  add  to  it  5  cc.  Na3Co(NO2)6  reagent,  and  let 
it  stand  for  10  minutes.     (Yellow  precipitate,  presence  of  POTASSIUM.) 
Notes. — 1.     This  confirmatory  test  should  not  be  omitted,  owing  to  the  pos- 
sibility that,  "in  consequence  of  imperfect  manipulation,  the  residue  left  undis- 
solved by  the  alcohol  in  P.  92  consists  of  NH4C1O4,  NaCl,  or  Na2SO4.     The 
NH4C1O4  may  result  from  incomplete  removal  of  the  ammonium  salts  in  the 
ignition,  the  NaCl  from  incomplete  conversion  of  the  chlorides  into  percMorates 
in  the  evaporation  with  HCIO^,  and  the  Na2SC>4  from  the  presence  of  sulfate 
which  was  not  removed  by  the  use  of  BaCl2. 

2.  The  boiling  with  NaNO2  and  HAc  serves  to  destroy  any  small  quantity 
of  ammonium  (up  to  20-30  mg.)  which  might  possibly  have  escaped  volatiliza- 
tion in  the  ignition.     It  is  destroyed  in  virtue  of  the  reaction  NH4+C1~+HNO2= 
N2+2H2O+H+C1~.     If  ammonium  is  present  in  the  solution   tested  with 
NasCo(NO2)6,  it  yields  a  precipitate  closely  resembling  that  produced  by  potas- 
sium. 

3.  By  this  test  the  presence  of  0.3  mg.  of  potassium  may  be  detected  within 
5—10  minutes.     The  yellow  color  of  the  precipitate  can  be  best  seen  by  collecting 
it  on  a  filter,  and  washing  out  the  solution  thoroly  with  cold  water. 

Procedure  94. — Detection  of  Sodium. — Pour  the  alcoholic  filtrate 
(P.  92)  into  a  dry  conical  flask  placed  in  a  vessel  of  cold  water,  and 
pass  into  it  a  fairly  rapid  current  of  dry  HC1  gas  until  the  gas  is  no 
longer  absorbed.  (White  precipitate,  presence  of  SODIUM.)  Treat 
the  precipitate  by  P.  95. 

Notes. — 1.  This  test  for  sodium  is  most  delicate  when  the  alcohol  is  com- 
pletely saturated  with  HC1  gas,  in  which  case  1  mg.  of  sodium  can  be  detected. 


P.  94  ANALYSIS  OF  THE  ALKALI-GROUP.  89 

2.  The  dry  HC1  gas  may  be  prepared  by  dropping  96%  H2SO4  from  a 
separating  funnel  into  a  flask  containing  solid  NaCl  covered  with  12-normal 
HC1,  and  passing  the  gas  through  a  gas- wash-bottle  containing  96%  HzSOi. 
Such  a  gas-generator  may  be  conveniently  kept  ready  for  use  in  a  hood  in  the 
laboratory,  as  the  evolution  of  gas  soon  ceases  when  no  more  H2SO4  is  added. 

3.  The  alcoholic  filtrate  containing  the  excess  of  HC1O4  must  not  be  evapo- 
rated, since  a  dangerous  explosion  is  likely  to  result. 

Procedure  95. — Confirmatory  Test  for  Sodium. — Filter  out  the  HC1 
precipitate  (P.  94),  wash  it  with  a  little  95%  alcohol,  pour  a  5-10  cc. 
portion  of  water  repeatedly  through  the  filter,  evaporate  the  solution 
just  to  dryness,  add  1  cc.  water  and  2  cc.  K2H2Sb2O7  reagent,  pour  the 
mixture  into  a  test-tube,  and  let  it  stand  for  at  least  half  an  hour,  pref- 
erably over  night.     (White  crystalline  precipitate,  presence  of  SODIUM.) 
Notes. — -1.     The  sodium  pyroantimonate  (Na2H2Sb2O7)  separates  as  a  heavy 
crystalline  precipitate,  which  usually  adheres  in  part  to  the  glass  in  the  form  of 
distinct  crystals,  which  can  be  best  seen  by  inverting  the  test-tube.     The  test 
is  not  extremely  delicate;  but  1  mg.  of  sodium  is  easily  detected. 

2.  With  the  antimonate  reagent  many  other  elements,  even  if  present  in 
small  quantity,  giv  precipitates;  thus  a  distinct  turbidity  is  produced  by  even 
0.1-0.2  mg.  of  calcium,  barium,  or  magnesium,  and  by  1-2  mg.  of  aluminum. 
These  elements  produce,  however,  light,  flocculent  precipitates  which  look  very 
different  from  the  heavy  crystalline  precipitate  obtained  with  sodium,  especially 
if  the  mixture  has  been  allowed  to  stand  a  few  hours.  The  crystals  of  the  so- 
dium salt  may  be  separated  from  a  flocculent  precipitate  by  shaking  the 
mixture,  waiting  long  enough  for  the  heavy  crystals  to  settle,  and  decanting 
off  the  solution  containing  the  flocculent  precipitate  in  suspension. 


90  SUPPLEMENTARY  PROCEDURES  FOR  BASES.  P.  97 

SUPPLEMENTARY   PROCEDURES   FOR  BASIC    CONSTITUENTS. 


Procedure  96. — Detection  of  Ammonium. — Place  0.2-0.3  g.  of  the 
finely  powdered  original  substance  and  2  cc.  NaOH  solution  in  a 
50  cc.  round-bottom  flask.  Insert  a  stopper  carrying  a  glass  rod 
around  whose  end  is  wound  a  piece  of  moist  red  litmus  paper;  and 
heat  the  mixture  nearly  to  boiling.  (Blue  coloration  of  the  litmus 
paper  and  odor  of  ammonia,  presence  of  AMMONIUM.) 

Confirmatory  Test  for  Ammonium. — If  the  litmus  turns  blue  or  the 
vapors  smell  of  ammonia,  pour  into  the  flask  10  cc.  water,  insert  a 
stopper  fitted  with  a  long,  wide  delivery  tube  leading  to  the  bottom 
of  a  test-tube  placed  in  a  vessel  of  cold  water,  and  distil  slowly  till 
about  half  the  water  has  passed  over.  To  the  distillate  add  a  10% 
solution  of  K2HgI4  in  3-normal  NaOH  drop  by  drop  so  long  as  the 
precipitate  increases.  (Orange  precipitate,  presence  of  AMMONIUM.) 

Notes. — 1.  Less  than  0.2  mg.  of  ammonium  can  be  detected  with  litmus 
paper  and  by  the  odor  when  the  test  is  carried  out  as  described  in  the  first 
paragraph  of  the  Procedure.  The  test  described  in  the  last  paragraph  is  use- 
ful with  very  small  quantities  of  ammonium  as  a  confirmation,  and  with  larger 
quantities  as  a  means  of  better  estimating  the  proportion  of  it  present. 

2.  The  orange  precipitate  produced  by  the  action  of  NH3  on  alkaline 
K2Hgl4  is  a  complex  compound  of  the  composition  HgO.HgNH2I.  The  test 
is  extremely  delicate,  a  distinct  precipitate  resulting  even  with  0.2  mg.  NH4 
in  5  cc.  solution,  and  a  pronounced  yellow  color  with  a  much  smaller  quan- 
tity. This  fact  must  be  taken  into  account  in  estimating  the  quantity  of 
ammonium  present. 

Procedure  97. — Determination  of  the  State  of  Oxidation  of  Mer- 
cury, Tin,  and  Iron. — If  mercury,  tin,  or  iron  has  been  found  to  be 
present,  heat  20-30  cc.  6-normal  H2S04  to  boiling  in  a  small  flask 
(to  expel  the  air  from  the  solution  and  flask);  drop  in  0.3  g.  of  the 
finely  powdered  original  substance;  and,  if  solution  does  not  take 
place  at  once,  boil  the  mixture  for  2-3  minutes,  covering  the  flask 
loosely  with  a  small  watch-glass.  Pour  5  cc.  portions  of  the  solution 
(through  a  filter  if  it  is  not  clear)  into  test-tubes  containing:  a,  5  cc. 
1-normal  HC1  (white  precipitate,  presence  of  MERCUROUS  MERCURY 
or  silver) ;  6,  5  cc.  SnCU  solution  (white  precipitate,  presence  of  MER- 
CURIC MERCURY)  ;  c,  5  cc.  HgCl2  solution  (white  precipitate,  presence 
of  STANNOUS  TIN);  d,  5  cc.  K3Fe(CN)6  solution  (blue  precipitate, 
presence  of  FERROUS  IRON)  ;  e}  5  cc.  KSCN  solution  (red  coloration, 
presence  of  FERRIC  IRON).  If  a  precipitate  is  produced  by  HC1  in  a, 
filter  and  add  the  filtrate  (instead  of  fresh  portions  of  the  H2S04  solu- 
tion) to  the  solutions  named  under  b  and  c;  and  if  silver  is  present 
in  the  substance,  treat  the  HC1  precipitate  with  NH4OH  by  P.  14. 


P.  97.  SUPPLEMENTARY  PROCEDURES  FOR  BASES.  91 

Notes. — 1.  If  in  preparing  the  solution  for  the  analysis  for  basic  constituents 
the  substance  was  dissolved  in  water  or  cold  dilute  HNO3,  the  state  of  oxidation 
of  mercury  will  have  been  determined  by  its  presence  or  absence  in  the  HC1 
and  H2S  precipitate.  But,  if  the  substance  was  treated  with  hot  or  concen- 
trated HNOa,  any  mercurous  compound  present  will  have  been  partly  or  com- 
pletely oxidized  to  the  mercuric  state. 

2.  Stannous  and  ferrous  salts  oxidize  rapidly  in  the  air.    Hence,  if  the  tests 
for  them  are  to  be  delicate,  the  contact  with  the  air  must  be  made  as  short  as 
possible. 

3.  Ferrous  and  ferric  salts  show  a  different  behavior  also  with  K4Fe(CN)e; 
the  former  giving  a  white  precipitate  (of  K2FeFe(CN)6)  which  rapidly  turns 
blue  in  the  air  and  the  latter  a  dark-blue  precipitate  (of  ferric  ferrocyanide, 
Fe^FeCeNe^).    Since  this  reagent  produces  a  precipitate  with  both  kinds  of 
salts,  it  is  less  suitable  than  K3Fe(CN)e  for  distinguishing  between  them. — 
The  precipitate  produced  by  K3Fe(CN)e  with  ferrous  salts  is  also  ferric  ferro- 
cyanide. 

Procedure  98. — Determination  of  the  State  of  Oxidation  of  Arsenic. — 
If  arsenic  has  been  found  to  be  present,  place  0.3  g.  of  the  solid 
substance  in  a  100  cc.  round-bottom  flask  arranged  for  distillation 
as  described  in  P.  101.  Place  in  the  receiving  flask  50  cc.  of  water. 
Pour  into  the  distilling  flask  10  cc.  12-normal  HC1  and  distil  till  about 
5  cc.  have  passed  over.  (See  Note  1.)  Pass  H2S  into  the  distillate. 
(Yellow  precipitate,  presence  of  ARSENITE.) 

If  H2S  produces  no  precipitate,  pour  into  the  distilling  flask  5  cc. 
12-normal  HC1  in  which  0.5  g.  powdered  FeS04  has  been  dissolved, 
and  distil  into  a  fresh  50  cc.  portion  of  water  till  5  cc.  have  passed 
over.  Pass  H2S  into  the  distillate.  (Yellow  precipitate,  presence  of 

ARSENATE.) 

Notes. — 1.  As  the  arsenic  vapors  are  very  poisonous,  care  must  be  taken  that 
they  do  not  escape  into  the  air. 

2.  This  method  of  distinguishing  arsenite  and  arsenate  depends  upon  the 
facts  that  arsenite  is  largely  converted  by  strong  HC1  into  AsCls,  which  is 
volatil  with  steam,  while  arsenate  is  not  converted  into  the  corresponding 
chloride,  and  is  therefore  not  volatil.     By  the  addition  of  FeSC>4,  however, 
arsenate  is  reduced  in  the  presence  of  HC1  to  AsCls,  which  then  volatilizes 
with  the  steam. — Less  than  1  mg.  of  arsenic  in  either  state  of  oxidation  is 
readily  detected  by  this  procedure. 

3.  When  arsenite  is  present,  it  can  be  completely  driven  over  into  the 
distillate  by  distilling  the  substance  with  10  cc.   12-normal  HC1  till  only 
2-3  cc.  remain,  replacing  the  HC1  which  has  distilled  off,  distilling  again,  and 
repeating    these  operations  till  the  distillate  givs  no  precipitate  with  B^S. 
Then,  to  test  for  arsenate,  a  solution  of  FeSC>4  in  HC1  may  be  added,  the 
mixture  again  distilled,  and  the  distillate  saturated  with  B^S.    The  removal 
of  the  arsenite,  is,  however,  so  slow  that  six  or  eight  repetitions  of  the  distilla- 
tion may  be  necessary. 


92  SUPPLEMENTARY  PROCEDURES  FOR  BASES.  P.  99 

Procedure  99. — Detection  of  Very  Small  Quantities  of  Arsenic  and 
Antimony. — Prepare  a  H2SO4  solution  of  the  substance  by  heating 
0.2-0.3  g.  of  it  with  10  cc.  H2SO4  if  it  does  not  contain  organic  matter, 
or  by  treating  a  larger  quantity  of  it  by  the  first  paragraph  of  P.  8 
if  it  contains  organic  matter.  Place  2-3  cc.  of  pure,  finely  granulated 
zinc  in  a  75-cc.  flask  fitted  with  a  stopper  through  which  pass  a 
thistle-tube  and  a  delivery-tube  bent  at  a  right  angle.  Connect  the 
delivery-tube  through  a  tube  filled  with  small  lumps  of  CaCl2  with  a 
hard-glass  tube  drawn  out  to  a  point  at  one  end  and  constricted  to  a 
fairly  wide  capillary  tube  in  the  middle.  Pour  into  the  flask  10-15 
cc.  water  and  10-12  drops  of  CuSO4  solution;  then  add  enough  H2S04 
to  produce  a  fairly  rapid  evolution  of  hydrogen.  After  the  air  has  been 
expelled  from  the  apparatus  (as  shown  by  the  fact  that,  on  filling  an 
inverted  7-cm.  test-tube  with  the  gas  and  touching  the  mouth  of  it  to  a 
gas-flame,  the  gas  burns  quietly  without  exploding)  light  the  gas  at 
the  end  of  the  hard-glass  tube  and  heat  that  tube  just  back  of  the  cap- 
illary with  a  small  gas  flame.  (See  Note  4.)  If  after  2-3  minutes  no 
black  deposit  appears  in  the  capillary  (showing  the  purity  of  the  re- 
agents), pour  into  the  flask  the  H2S04  solution  of  the  substance,  a  little 
at  a  time,  and  let  the  gas  continue  to  pass  through  the  heated  tube  for 
4-5  minutes.  (Black  deposit  in  the  capillary,  presence  of  ARSENIC 
or  ANTIMONY.)  Cool  the  hard-glass  tube,  and  dip  the  capillary  part 
of  it  in  a  test-tube  containing  NaOCl  solution.  (Partial  or  complete 
solution  of  the  deposit,  presence  of  ARSENIC;  incomplete  solution  of 
the  deposit,  presence  of  ANTIMONY.) 

Notes. — 1.  This  method  of  detecting  arsenic  and  antimony  depends  upon  the 
facts  that,  when  hydrogen  is  produced  in  a  solution  containing  these  elements 
all  of  the  arsenic  is  converted  into  hydrogen  arsenide  gas  (AsHs)  and  a  part  of 
the  antimony  is  converted  into  hydrogen  antimonide  gas  (SbHs) ;  and  that  these 
gases  decompose  into  their  elements  at  a  moderately  high  temperature. 

2.  The  treatment  of  the  deposit  with  NaOCl  solution  serves  to  distinguish  the 
two  elements;  for  metallic  arsenic  is  readily  converted  into  HsAsO4  by  that 
reagent,  while  metallic  antimony  is  not  acted  upon  by  it.     Even  when  they  are 
present  together,  it  is  usually  possible  to  detect  both  of  them;  for  arsenic,  being 
more  volatil,  deposits  in  the  part  of  the  capillary  further  from  the  flame,  and  this 
part  of  the  deposit  may  be  seen  to  dissolve,  and  the  other  part  to  remain  undis- 
solved,  in  the  treatment  with  NaOCl  solution. 

3.  The  addition  of  the  CuSO4  solution  serves  to  produce  a  deposit  of  copper 
on  the  zinc  granules  and  thus  to  accelerate  through  the  voltaic  action  the  evo- 
lution of  hydrogen  and  the  formation  of  the  AsHs  and  SbHs. 

4.  Hydrogen  arsenide  is  an  extremely  poisonous  gas.     Therefore  it  must  not 
be  allowed  to  escape  into  the  air  through  leaks  in  the  apparatus  or  by  discon- 
tinuing the  heating  of  the  hard-glass  tube  while  much  AsH3  is  being  evolved. 

5.  This  process  is  commonly  employed  for  the  detection  of  arsenic  in  papers, 
fabrics,  and  other  organic  materials. 


DETECTION  OF  THE  ACIDIC  CONSTITUENTS. 


GENERAL   DISCUSSION. 


The  acidic  constituents  whose  detection  is  here  provided  for  are : 

Carbonate.  Cyanide.  Bromate. 

Sulfite.  Chloride.  Nitrate. 

Thiosulfate.  Bromide.  Phosphate. 

Nitrite.  Iodide.  Sulfate. 

Hypochlorite.  Thiocyanate.  Borate. 

Sulfide.  Chlorate.  Fluoride. 

Somewhat  different  processes  are  described  in  this  book  for  the  de- 
tection of  these  constituents,  according  as  the  substance  is: 

(1)  completely  dissolved  by  cold  dilute  acids; 

(2)  decomposed  only  by  hot  concentrated  acids; 

(3)  not  decomposed  even  by  concentrated  acids. 

In  the  process  employed  when  the  substance  is  dissolved  by  cold 
dilute  acids,  the  constituents  that  yield  with  acids  readily  volatil  prod- 
ucts (hereafter  called  for  short  the  readily  volatil  constituents),  namely, 
carbonate,  sulfite,  thiosulfate,  sulfide,  nitrite,  hypochlorite,  and  cyan- 
ide, are  first  tested  for  by  treating  the  substance  with  hot  dilute  sulf uric 
acid  and  bringing  suitable  solutions  in  the  form  of  drops  or  test- 
papers  into  contact  with  the  vapors.  This  process  is  outlined  in  Table 

XIII  on  page  96.     A  nitric  acid  solution  of  the  substance  is  then  pre- 
pared; and,  by  adding  to  portions  of  it  various  reagents,  most  of  the 
other  acidic  constituents  are  detected  through  the  formation  of  char- 
acteristic precipitates  or  colorations.     This  process  is  outlined  in  Tables 

XIV  and  XV  on  pages  97  and  100.     The  remaining  constituents  are 
finally  tested  for  with  samples  of  the  original  substance,  as  outlined 
in  Table  XIX  on  page  111. 

In  the  process  employed  when  the  substance  is  decomposed  only  by 
hot  concentrated  acids  the  substance  is  first  boiled  with  a  solution  of 
phosphoric  acid.  The  vapors  which  come  off  while  the  phosphoric 
acid  is  dilute  contain  the  readily  volatil  constituents;  these  are  con- 
densed in  barium  hydroxide  solution,  constituting  the  " first  distillate." 
The  vapors  passing  over  when  the  phosphoric  acid  becomes  more 
concentrated  contain  the  acids  corresponding  to  the  less  volatil  con- 
stituents, namely,  to  chloride,  bromide,  iodide,  thiocyanate,  chlorate, 

93 


94  DETECTION  OF  ACIDIC  CONSTITUENTS. 

and  nitrate;  these  are  condensed  in  water,  constituting  the  "second 
distillate."  Finally  when  the  phosphoric  acid  has  become  very  con- 
centrated, copper  is  added  to  it,  causing  sulfate  to  be  reduced  to  sulfite 
and  to  pass  over  as  sulfurous  acid  into  the  "third  distillate."  These 
distillates  are  then  tested  for  the  separate  constituents  by  appropriate 
reagents.  This  system  of  procedure  is  summarized  in  Tables  XVII 
and  XVIII  on  pages  105  and  109.  Portions  of  the  solid  substance  are 
then  tested  for  the  remaining  constituents  by  such  of  the  procedures 
outlined  in  Table  XIX  on  page  111  as  the  results  of  the  previous  tests 
make  necessary. 

The  second  of  these  processes  can,  of  course,  be  employed,  in  place 
of  the  first  one,  also  with  substances  which  are  dissolved  by  cold  dilute 
acids;  and  the  analyst  may  prefer  to  use  it  as  the  general  procedure  for 
all  substances  decomposable  by  acids. 

In  the  process  employed  with  substances  not  completely  decomposed 
even  by  hot  concentrated  acids,  samples  of  the  solid  substance  are 
tested  for  the  readily  volatil  constituents  and  portions  of  a  nitric  acid 
extract  of  the  substance  are  tested  for  the  other  constituents,  just  as 
in  the  case  of  substances  dissolved  completely  by  dilute  nitric  acid. 
The  residue  undissolved  by  nitric  acid  is  then  fused  with  sodium  car- 
bonate, the  mass  is  extracted  with  water,  and  the  solution  is  tested  for 
the  constituents  that  are  likely  to  be  present  in  insoluble  substances, 
namely,  for  sulfate,  sulfide,  chloride,  phosphate,  borate,  and  fluoride. 

It  is  to  be  noted  that  the  system  of  procedure  for  detecting  the  acidic 
constituents  can  often  be  much  shortened  by  omitting  the  tests  for 
certain  constituents  which  are  excluded  by  the  known  character  or 
source  of  the  substance,  or  by  its  solubility  considered  in  connection 
with  the  basic  constituents  present.  Thus,  it  is  useless  to  test  a  mineral 
for  nitrite,  sulfite,  hypochlorite,  chlorate,  or  cyanide;  or,  in  a  neutral 
water-soluble  substance  containing  barium  or  silver  it  is  unnecessary 
to  test  for  any  of  the  acidic  constituents  which  form  insoluble  com- 
pounds with  these  elements.  A  general  statement  as  to  the  solubilities 
of  substances  in  water  or  dilute  acids  will  be  found  in  Note  9  on  page 
31;  and  numerical  values  of  the  solubilities  of  some  analytically  im- 
portant substances  are  given  in  a  table  on  page  124. 

In  addition  to  the  acidic  constituents  listed  on  page  93,  the  follow- 
ing ones,  which  are  detected  in  the  course  of  the  analysis  for  basic  con- 
stituents, are  frequently  met  with  in  minerals  or  industrial  products : 

Silicate.  Arsenate.  Chromate. 

Stannate.  Areenite.  Permanganate. 


P.  100  DETECTION  OF  ACIDIC  CONSTITUENTS.  95 

GENERAL   DIRECTIONS. 


Procedure  100. — General  Directions. — If  the  substance  is  completely 
dissolved  by  cold  dilute  HNOs  (as  used  in  P.  2),  test  samples  of  the 
substance  for  readily  volatil  constituents  by  P.  101.  Prepare  a  solu- 
tion of  the  substance  by  dissolving  2  g.  of  it  in  30  cc.  1-normal  HN03; 
test  portions  of  this  solution  by  P.  102-104;  and  test  the  remainder  of 
the  solution  by  P.  105  and  P.  106,  if  in  P.  103  halides  are  found  to  be 
present.  Test  fresh  samples  of  the  substance  for  borate  by  P.  121 
and  for  nitrate  by  P.  124;  also  for  nitrite  by  P.  125  and  for  hypochlorite 
by  P.  126,  if  the  previous  tests  show  that  they  may  be  present. 

If  the  substance  is  not  completely  dissolved  by  cold  dilute  HN03, 
but  is  decomposed  by  hot  concentrated  acids  (as  used  in  P.  3  and  4), 
treat  a  2  g.  portion  of  it  by  P.  Ill,  and  treat  the  distillates  thus  ob- 
tained as  directed  in  P.  111.  Test  also  fresh  samples  of  the  sub- 
stance for  borate,  fluoride,  and  phosphate  by  P.  121,  122,  and  123; 
also  for  nitrite,  hypochlorite,  and  chlorate  by  P.  124-127,  if  the 
previous  tests  show  that  these  constituents  may  be  present. 

If  the  substance  is  not  completely  decomposed  even  by  concentrated 
acids,  treat  samples  of  the  substance  by  P.  131. 

Notes. — 1.  When  the  substance  is  completely  decomposed  by  dilute  acids 
the  analysis  for  acidic  constituents  can  be  made  by  testing  for  the  various 
constituents  in  the  manner  described  in  the  first  paragraph  of  the  foregoing 
Procedure.  When,  however,  the  substance  is  not  decomposed  by  dilute  acids, 
some  of  the  acidic  constituents  might  escape  detection  if  tested  for  in  these 
ways.  With  such  a  substance  it  is  therefore  necessary  to  use  a  more  powerful 
decomposing  agent.  Phosphoric  acid  is  suitable  for  this  purpose,  since  it  is  a 
fairly  strong  acid  whose  solution  can  be  made  highly  concentrated  without  caus- 
ing much  volatilization  or  decomposition  of  the  acid.  When  the  substance 
does  not  dissolve  completely  in  dilute  HNOs  it  is  therefore  directed  to  boil  it 
with  dilute  HsPCU  (as  described  in  P.  Ill)  and  to  condense  the  vapors  which 
contain  all  the  volatil  acids  resulting  from  the  decomposition  of  the  substance. 
This  process  serves  at  the  same  time  to  separate  the  acidic  constituents  of  the 
substance  from  its  basic  constituents,  thereby  facilitating  the  detection  of  the 
acidic  constituents. 

2.  Even  with  a  substance  which  dissolves  completely  in  dilute  acids  it  is 
often  advantageous  to  employ  the  distillation  process  of  P.  Ill;  for  tho  the 
tests  for  the  readily  volatil  acidic  constituents  described  in  P.  101  are  very 
delicate,  they  do  not  enable  a  satisfactory  estimate  to  be  made  of  the  propor- 
tions in  which  the  constituents  are  present.     Moreover,  these  tests  fail  to  detect 
certain  of  the  constituents  when  they  are  present  together  (for  example,  car- 
bonate in  the  presence  of  sulfite). 

3.  If  the  substance  is  not  decomposed  even  by  hot  concentrated  acids,  it  has 
to  be  decomposed  by  fusion  with  Na^OOa,  as  described  in  P.  131. 


96 


DETECTION  OF  ACIDIC  CONSTITUENTS. 


P.  101 


SUBSTANCES    DISSOLVED    BY   DILUTE   NITRIC    ACID. 


TABLE  XIII. — DETECTION  OF  THE  READILY  VOLATIL  ACIDIC  CONSTITUENTS. 


Heat  the  substance  with  dilute  #2£04  (P.  101). 


Vapors:  CO2,  SO2,  H2S,  NO2,  C12,  Br2,  I2,  HCN.    Expose  to  the  vapors: 


Ba(OH}z  solution. 

PbAc%  paper. 

Starch  and  KI 
paper. 

Fe(OH)2,  Fe(OH^, 
and  NaOH  on  paper. 

White  turbidity: 
BaCO3  or  BaSO3. 

(ShoWS  CARBONATE, 
STJLFITE,  01  THIO- 

SULFATE.) 

Black  color: 
PbS. 

(ShoWS  STTLFIDE.) 

Blue  color:  I2. 
(Shows  NITRITE, 

HYPOCHLORITE, 
CHLORATE,  BRO- 
MATE,  Or  IODIDE.) 

Formation  of 
Na4Fe(CN)6. 
Dip  in  HCl. 

Blue  color: 

Fe4(Fe(CN)6)3. 

(ShoWS  CYANIDE.) 

Procedure  101. — Detection  of  Constituents  Yielding  Readily  Volatil 
Products. — Place  0.3  g.  of  the  finely  powdered  substance  in  a  30  cc. 
conical  flask,  add  2  cc.  of  water,  heat  the  mixture  nearly  to  boiling, 
and  to  the  hot  mixture  add  5  or  6  drops  H2S04.  Note  whether  there 
is  an  odor  or  formation  of  gas  bubbles.  Insert  in  the  flask  a  two-hole 
rubber  stopper  through  which  passes  a  glass  rod  from  which  is  sus- 
pended a  drop  of  Ba(OH)2  solution.  (White  precipitate,  presence 
of  CARBONATE,  SULFITE,  or  THiosuLFATE.)  Make  B,  conical  paper 
roll  out  of  half  a  filter-paper;  insert  the  narrow  end  of  it,  in  place 
of  the  glass  rod,  in  one  hole  of  the  rubber  stopper;  dip  the  other 
end  of  the  roll  in  PbAc2  solution  contained  in  a  7-cm.  test-tube; 
insert  the  stopper  in  the  flask;  and  heat  the  liquid  nearly  to  boiling. 
(Blackening  of  the  paper,  presence  of  SULFIDE.)  Replace  the  paper 
roll  by  a  fresh  one  which  has  been  dipped  in  a  solution  of  KI  and 
starch;  and  heat  the  liquid  again  nearly  to  boiling.  (Blue  coloration 
of  the  paper,  presence  of  NITRITE,  HYPOCHLORITE,  CHLORATE,  BRO- 
MATE,  or  IODIDE.) — Place  in  the  flask  a  fresh  0.3  g.  sample  of  the  sub- 
stance, and  add  to  it  2  cc.  of  water  and  5-6  drops  of  H2SO4;  insert  a 
rubber  stopper  carrying  a  paper-roll  which  has  been  dipped  first  in  a 
2-normal  FeS04  solution  and  then  in  a  1-normal  NaOH  solution;  and 
heat  the  mixture  in  the  flask  nearly  to  boiling.  Remove  the  roll,  and 
dip  it  first  in  HCl  and  then  in  water.  (Blue  coloration  of  the  paper, 
presence  of  CYANIDE.) 


P.  101 


DETECTION  OF  ACIDIC  CONSTITUENTS. 


97 


Notes. — 1.  These  tests  are  all  delicate  enough  to  show  0.1-0.2  mg.  of 
the  respectiv  constituents.  When  any  of  the  tests  yields  a  positiv  result, 
the  nature  and  quantity  of  the  constituent  in  the  substance  which  givs  rise  to 
the  test  is  more  definitly  determind  by  later  procedures,  as  described  in  the 
general  directions  given  in  P.  100. 

2.  A  blue  coloration  of  the  starch-KI  paper  shows  that  there  is  present  in 
the  vapors  either  free  iodin  or  one  of  the  volatil  substances  which  liberates  iodin 
from  KI;  namely,  chlorin,  bromin,  or  nitrogen  dioxide.  In  the  presence  of  starch, 
which  forms  with  water  a  colloidal  solution,  the  iodin  is  dissolved  by  the  minute 
starch  globules  or  is  adsorbed  on  their  surface;  and  in  this  finely  divided  state 
it  possesses  a  deep  blue  color.  Chlorin  or  bromin  usually  arises  from  the  presence 
in  the  substance  of  hypochlorite,  chlorate,  or  bromate,  together  with  a  chloride 
or  bromide.  Iodin  may  be  liberated  from  an  iodide,  when  an  oxidizing  substance, 
such  as  a  ferric  salt,  is  also  present. 


TABLE  XIV. — DETECTION  OF  THE  ACIDIC  CONSTITUENTS  PRECIPITATED  FROM  ACID 
SOLUTIONS  BY  BARIUM  AND  SILVER  SALTS. 


To  a  HNOz  solution  of  the  substance 
add  BaCk  (P.  102). 

To  a  HNOs  solution  of  the  substance 
add  Cd(N03)2  (P.  10S). 

Precipi- 
tate: 
BaSO4. 
(Shows 

SUL- 
FATE.) 

Filtrate.    Add  Br2. 

Yellow 
precipitate: 
CdS. 
(Shows 

SULFIDE.) 

Filtrate:  add  AgN03. 

Precipi- 
tate: 
BaSO4. 

(Shows 

SULFITE.) 

Filtrate.  Add  NH*  Ac. 

Precipitate: 
AgCl, 
AgBr,  Agl, 

Ag2(CN)2, 
AgSCN. 
(Shows 

HALIDES, 
CYANIDE, 

or  THIO- 

CYANATE.) 

Filtrate: 
AgC103, 
AgBr03. 
Add  H2S03. 

Yellow 
precipitate: 
BaCr04. 
(Shows 

CHROMATE.) 

Filtrate. 
Add  CaCk. 

Precipitate: 
CaF2. 

(Shows 

FLUORIDE.) 

Precipitate: 
AgCl,  AgBr. 
(Shows 

CHLORATE 
Or  BROMATE.) 

Procedure  102. — Detection  of  Sulfate,  Sulfite,  Chr ornate,  and  Flu- 
oride.—To  10  cc.  of  the  HNO3  solution  (P.  100)  add  10  cc.  BaCl2  solu- 
tion, and  let  the  mixture  stand  for  5  minutes.  (White  precipitate, 
presence  of  SULFATE.) 

Filter  the  mixture,  repeatedly  if  necessary;  add  to  the  filtrate  (unless 
it  smells  of  H2S)  saturated  Br2  solution,  1  cc.  at  a  time,  till  it  is  present 
in  excess,  and  let  the  mixture  stand  for  5  minutes.  (White  precipitate, 
presence  of  SULFITE.) 

Filter  the  mixture,  repeatedly  if  necessary,  add  to  the  filtrate  5  cc. 
3-normal  NH4Ac  solution,  and  let  the  mixture  stand  5  minutes.  (Fine 
yellow  precipitate,  presence  of  CHROMATE.) 


98  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  102 

Filter  the  mixture,  add  to  the  filtrate  10  cc.  CaCl2  solution,  and  let 
it  stand  15  minutes.  (White  turbidity,  presence  of  FLUORIDE.)  Confirm 
the  presence  of  fluoride  by  filtering  off  the  precipitate,  igniting  the 
filter  containing  it  in  a  spiral  of  platinum  wire  till  it  is  incinerated, 
and  treating  the  residue  by  P.  122. 

Notes. — 1.  Of  the  barium  salts  the  sulfate  is  the  only  one  that  is  precipitated 
by  a  moderate  excess  of  BaC^  from  a  HNOs  solution  as  strong  as  0.5  normal. 
From  the  HAc  solution  produced  by  adding  to  such  a  HNOs  solution  1?  times 
as  many  equivalents  of  NH4Ac  the  chromate  is  precipitated  completely,  and  the 
fluoride  is  precipitated  when  more  than  about  10  mg.  of  fluorin  is  present. 
Sulfite,  if  it  were  not  previously  removed  by  converting  it  into  sulfate  by  the 
addition  of  Br2,  would  also  precipitate  on  the  addition  of  the  NH4Ac.  From 
a  neutral  solution  phosphate,  carbonate,  and  borate  are  also  precipitated;  but 
none  of  these  separates  from  the  HAc  solution. 

2.  Of  all  the  inorganic  acidic  constituents  the  fluoride  is  the  only  one  whose 
calcium  salt  is  much  less  soluble  in  water  than  its  barium  salts.     BaF2  is  soluble 
in  water  to  such  an  extent  that  under  the  conditions  of  the  procedure  about  10 
mg.  of  fluorin  remain  in  the  HAc  solution.     A  precipitate  produced  by  CaCfe 
therefore  shows  fluoride.     The  precipitate  of  CaF2  has,  moreover,  a  characteristic 
appearance,  separating  first  as  a  milky  turbidity  which  slowly  settles  out  in 
flocculent  form. 

3.  The  addition  of  NH*Ac  to  the  HNOs  solution  may  cause  the  precipitation 
of  other  substances  than  BaCrO4  and  BaF2;  namely,  of  any  substances,  such  as 
bismuth  salts,  ferric  or  aluminum  phosphate,  silicic  acid,  which  are  dissolved  by 
HNOs  hut  not  by  HAc.     In  such  a  case  the  test  for  chromate  is  obscured;  but 
fluoride  may  still  be  detected  in  the  nitrate  with  CaCfe. 

4.  Sulfide  and  sulfite  can  not  be  present  together  in  an  acid  solution;  for  they 
destroy  each  other  with  the  separation  of  sulfur.     It  is  therefore  useless  to  test 
a  solution  containing  H2S  for  sulfite;  moreover,  the  sulfur  which  would  be  precip- 
itated on  adding  bromin  to  such  a  solution  might  be  mistaken  for  BaSO4. 

Procedure  103. — Detection  of  Sulfide,  of  Halides,  and  of  Chlorate  or 
Bromate. — To  a  5  cc.  portion  of  the  HNOs  solution  (P.  100)  add  5  cc. 
Cd(NO3)2  solution.  (Yellow  precipitate,  presence  of  SULFIDE.) 

Filter  out  the  precipitate;  and  to  the  filtrate  add  20  cc.  of  water  and 
5  cc.  AgNOs  solution.  (White  precipitate,  presence  of  CHLORIDE, 
CYANIDE,  or  THIOCYANATE;  yellow  precipitate,  presence  of  BROMIDE  or 
IODIDE.) 

Filter  out  the  precipitate;  to  the  filtrate  add  a  few  drops  of  AgNOs 
solution;  then  add  20  cc.  HNOs  and  5  cc.  saturated  S02  solution,  and 
let  the  mixture  stand  for  5  minutes.  If  there  is  a  precipitate,  heat  the 
mixture  nearly  to  boiling.  (White  precipitate,  presence  of  CHLORATE; 
yellow  precipitate,  presence  of  BROMATE.) 


P.  108  DETECTION  OF  ACIDIC  CONSTITUENTS.  99 

Notes. — 1.     The  presence  of  sulfide  is  detected  by  the  test-paper  test  in  P. 
101;  but  its  precipitation  as  CdS  enables  the  amount  to  be  better  estimated. 

2.  All  the  common  silver  salts,  except  the  three  halides,  the  cyanide,  thio- 
cyanate,  and  sulfide,  are  either  soluble  in  water  (as  are  the  nitrate,  sulfate, 
chlorate,  and  fluoride),  or  dissolve  readily  in  HNOs  owing  to  displacement  of 
the  weaker  acid  (as  do  the  phosphate,  carbonate,  borate,  and  sulfite).    There  are, 
however,  some  exceptions  to  the  principle  that  salts  of  weak  acids  are  readily 
soluble  in  a  strong  acid.    Thus  A&S  does  not  dissolve  in  dilute  HNO3  because 
its  solubility  in  pure  water  is  so  extremely  small  that  there  is  only  a  very  minute 
concentration  of  S~  ion  in  the  saturated  solution,  and  this  can  yield,  in  ac- 
cordance with  the  mass-action  law,  only  a  relativly  small  concentration  of 
SH~  and  unionized  H2S  with  the  H+  ion  of  the  HNOs.    Silver  cyanide  has 
for  another  reason  a  very  slight  concentration  of  its  anion  in  its  saturated 
solution;  namely,  because  of  the  fact  that  this  salt  exists  in  the  solution  mainly 
as  Ag+  and  Ag(CN)2~,  and  scarcely  at  all  as  Ag+  and  CN~. 

3.  The  reduction  of  chlorate  to  chloride  by  the  H2SOs  is  not  instantaneous; 
but  it  is  so  rapid  that  0.5  mg.  ClOs  in  a  volume  of  60  cc.  produces  a  precipitate 
in  less  than  5  minutes. 

4.  Before  the  addition  of  the  H2SOs  a  few  drops  of  AgNOs  are  added  to  make 
sure  that  the  halides  have  been  completely  precipitated.     A  large  quantity  of 
HNOs  is  also  added  to  prevent  the  precipitation  of  A^SOs;  and,  if  a  precipitate 
separates  on  adding  the  SO2  solution,  the  mixture  is  heated  nearly  to  boiling  to 
make  sure  that  the  precipitate  is  not  Ag^Os. 

5.  If  much  bromate  is  present,  some  of  it  is  precipitated  upon  the  first 
addition  of  AgNOs,  along  with  the  silver  halides;    but  some  of  it  also  remains 
in  the  solution  and  shows  the  same  behavior  as  chlorate.    In  order  to  distinguish 
between  them,  the  final  precipitate  with  AgNOs  may  be  treated  as  follows: 
Suspend  it  in  25  cc.  water,  pass  in  H2S  until  the  mixture  is  saturated  with  it, 
heat  to  boiling,  filter  off  the  precipitated  Ag2S,  boil  the  filtrate  till  the  H2S  is 
expelled,  and  test  it  for  bromide  and  chloride  by  P.  106. 


100 


DETECTION  OF  ACIDIC  CONSTITUENTS. 


P.  104 


TABLE  XV. — DETECTION  OF  PHOSPHATE  AND  THE  SEPARATE  HALIDES. 
To  portions  of  the  HNOS  solution  of  the  substance  (P.  100} : 


Add  (NHAzMoOi 
(P.  104). 

Add  FeCk 
(P.  105)* 

Add  NaAc,  HAc,  KMnO*,  and  CHCk  (P.  106)* 

Chloroform 
layer,  purple: 
I2. 

(ShoWS  IODIDE.) 

Water  layer:  add  HzSO*,  more 
KMnOi  and  CHCk. 

Yellow  precipitate: 
(NH4)3P04. 
12  MoO3. 

(ShoWS  PHOS- 
PHATE.) 

Red  color: 

Fe(SCN)3. 
(Shows  THIO- 

CYANATE.) 

Chloroform 
layer,  orange: 
Br2. 

(ShoWS  BRO- 
MIDE.) 

Water  layer: 
boil  out  the  Brz, 
add  HNO&   and 
AgNOz. 

Precipitate: 
AgCl 

(ShoWS  CHLO- 
RIDE.) 

•These  procedures  are  followed  only  in  case  AgNOj  produces  a  precipitate  in  P.  103. 

Procedure  104. — Detection  of  Phosphate. — Add  3  cc.  of  the  HN03 
solution  (P.  100)  to  6-8  cc.  (NH^MoC^  solution,  and  let  the  mixture 
stand  5-10  minutes.  (Yellow  precipitate,  presence  of  PHOSPHATE.) 

Notes. — 1.  The  yellow  precipitate  produced  is  a  complex  compound,  am- 
monium phosphomolybdate,  of  the  composition  (NH4)3PO4.12MoOa 

2.  In  order  that  the  test  may  be  delicate,  a  large  proportion  of  the 
(NEU)2MoO4  must  be  present  to  reduce  the  solubility  of  the  precipitate;  and 
a  short  time  must  be  allowed  for  the  formation  of  the  complex  phosphomolyb- 
date. This  is  promoted  by  gentle  warming;  but  in  a  hot  solution  arsenate  or 
silicate  may  giv  rise  to  a  similar  yellow  precipitate,  while  in  the  cold  the  re- 
action is  given  only  by  phosphate.  By  this  test  0.1  mg.  P(>4  may  be  easily 
detected.  The  great  delicacy  of  this  test  should  be  borne  in  mind  in  estimating 
the  quantity  of  phosphate  present. 

Procedure  105. — Detection  of  Thiocyanate. — If  AgN03  produced 
a  precipitate  (in  P.  103),  add  to  2  cc.  of  the  HNO3  solution  (P.  100) 
4-5  drops  of  FeCl3  solution.  (Red  color,  presence  of  THIOCYANATE.) 

Note. — The  red  color  arises  from  the  formation  by  metathesis  of  Fe(SCN)3,  a 
substance  whose  degree  of  ionization  is  relativly  small.  A  distinct  reddish- 
yellow  color  is  produced  by  0.1  mg.  SCN,  a  deep-red  color  by  1  mg.  or  more. 

Procedure  106. — Detection  of  the  Separate  Halides. — If  AgNOa  pro- 
duced a  precipitate  (in  P.  103),  add  to  the  remaining  10  cc.  of  the 
HN03  solution  (P.  100)  in  a  conical  flask  3-normal  Na2CO3  solu- 
tion, a  few  drops  at  a  time,  till  the  liquid  no  longer  givs  a  decided 
red  color  to  blue  litmus-paper.  (If  too  much  has  been  accidentally 
added,  add  HN03  drop  by  drop  till  the  solution  again  reddens  blue 


P.  106  DETECTION  OF  ACIDIC!  pQ%J[TfSpjJXyfai'  101 

litmus-paper.)  Then  add  8  cc.  NaAc  solution,  2  cc.  HAc,  and  (after 
filtering  out  any  precipitate)  3  cc.  chloroform  (CHC13).  Finally  add 
1%  KMn04  solution,  1  cc.  at  a  time,  shaking  vigorously  after  each 
addition,  till  the  aqueous  layer  becomes  pink.  (Purple  coloration  of 
the  chloroform,  presence  of  IODIDE.)  Pour  the  mixture  through  a 
moistened  filter  to  remove  the  chloroform  and  precipitated  Mn02,  and 
shake  the  filtrate  once  or  twice  with  a  fresh  10  cc.  portion  of  chloroform 
to  extract  all  the  iodin. 

Place  the  aqueous  solution  and  3  cc.  chloroform  in  a  separating 
funnel,  add  5  cc.  H2S04,  and  1  cc.  1%  KMn04  solution,  unless  such 
an  excess  is  already  present,  and  shake.  (Yellow  or  orange  coloration 
of  the  chloroform,  presence  of  BROMIDE.) 

Transfer  the  aqueous  layer  to  a  casserole,  add  5-20  cc.  1%  KMn04 
solution,  and  boil  the  mixture  3-5  minutes,  or  until  the  volume  has 
been  reduced  to  10  cc.  Filter  off  the  Mn02,  and,  if  the  solution  is 
still  pink,  add  H2SOs  solution  drop  by  drop  until  it  is  colorless.  Dilute 
the  solution  to  100  cc.,  filter  if  necessary,  and  add  20  cc.  HN03  and 
5  cc.  AgNOs  solution.  (White  precipitate,  presence  of  CHLORIDE.) 

Notes. — 1.  This  separation  is  based  upon  the  different  rates  at  which 
KMnO4  sets  free  by  oxidation  the  three  halogens  from  their  salts  in  a  solution 
of  definit  hydrogen-ion  (H+)  concentration.  A  dilute  solution  of  acetic  acid 
containing  considerable  sodium  acetate  has  such  a  hydrogen-ion  concentration 
that  an  iodide  is  immediately  oxidized  by  KMnO4  with  liberation  of  iodin, 
while  bromide  and  chloride  are  not  oxidized  to  an  appreciable  extent  in  the 
time  required  for  the  operations.  When  the  H+  concentration  is  increased  by 
the  addition  of  the  prescribed  quantity  of  H2S04,  the  bromide  is  oxidized  very 
rapidly,  while  the  rate  of  the  corresponding  reaction  for  the  chloride  is  still  so 
small  at  room  temperature  that  scarcely  any  chlorin  is  set  free.  Even  when 
the  solution  is  boiled  to  expel  the  bromin,  only  a  small  fraction  of  the  chloride 
present  is  oxidized  to  chlorin. 

2.  To  secure  satisfactory  results,  the  directions  as  to  the  quantities  of  the 
acids  added  must  be  followed  carefully.    The  proper  quantity  of  H2SO4  is  that 
required  to  react  with  all  the  sodium  acetate  and  to  giv  in  addition  an  excess 
equal  to  about  1  cc.  H2SO4  per  20  cc.  of  solution. 

3.  The  yellow  color  of  bromin  in  3-5  cc.  chloroform  enables  about  0.5  mg. 
Br  to  be  detected  in  this  procedure. 

4.  A  very  small  precipitate  of  AgCl  obtained  at  the  end  of  the  procedure 
does  not  necessarily  indicate  the  presence  of  chloride  in  the  substance,  unless 
the  reagents  used  have  been  proved  to  be  entirely  free  from  chloride.    Even 
then  a  very  slight  precipitate  (corresponding  to  less  than  0.1  mg.  Cl)  may 
result  from  a  reaction  between  the  permanganate  and  chloroform.    For  these 
reasons  a  blank  test  should  be  made  in  any  doubtful  case. 

5.  Before  adding  the  AgNOs  in  the  test  for  chloride  the  solution  is  diluted 
and  HN03  is  added  to  it,  so  as  to  prevent  the  precipitation  of  A&SC^  and  Ag2SOs. 

6.  If  HCN,  H2S,  or  HSCN  is  present  in  the  solution,  it  will  be  expelled 
or  destroyed  by  the  boiling  with  KMnO4  before  the  test  for  chloride  is  made. 


102  '.ti 


#F  ACIDIC  CONSTITUENTS. 


P.  Ill 


SUBSTANCES   DECOMPOSED    ONLY   BY    CONCENTRATED   ACIDS. 


TABLE  XVI. — BEHAVIOR  OF  THE  ACIDIC  CONSTITUENTS  ON  DISTILLATION  WITH 

PHOSPHORIC  Aero. 

Distil  the  substance  with  dilute  H^PO*  (P.  111).  Collect  the  first  half  of  the  distillate 
in  Ba(OH)z  solution  and  the  second  half  in  water.  To  the  residue  add  Cu  and  distil 
again,  collecting  this  third  distillate  in  water. 


FIRST  DISTILLATE 

SECOND   DISTILLATE 

THIRD  DIST. 

NONVOLATIL  RESIDUE 

COo  from  carbonate. 

HC1  from  chloride. 

SO2  from 

HPOa  from  phosphate. 

862  from  sulfite  or 

HBr  from  bromide. 

sulfate. 

HBO2  from  borate. 

thiosulfate. 

HI  from  iodide. 

H2SiOs  from  silicate. 

Cl2  from  hypochlorite, 

HSCN  from  thio- 

chlorate,  or  chloride.* 

cyanate. 

Bis  from  bromate  or 

HCN  from  f  erro-  or 

bromide.* 

ferri-cyanide. 

Ij  from  iodide.* 

H2S  from  insoluble 

HNO2  from  nitrite. 

sulfides. 

H2S  from  sulfide. 

HNOs  from  nitrate. 

HCN  from  cyanide. 

Ck  from  chlorate  or 

chloride.* 

Br2  from  bromide.* 

12  from  iodide. 

*  When  the  substance  contains  also  an  oxidizing  compound. 

Procedure  in. — Distillation  with  Phosphoric  Add. — Place  2  g.  of 
the  finely  powdered  substance  and  a  few  glass  beads  in  a  distilla- 
tion-apparatus, arranged  as  shown  in  the  figure,  consisting  of  a  100  cc. 
round-bottom  hard-glass  flask  fitted  with  a  rubber  stopper,  through 
which  pass  a  delivery  tube  and  a  safety-tube,  20-30  cm.  long,  leading 
to  the  bottom  of  the  flask.  Support  the  flask  in  an  inclined  posi- 
tion. Lead  the  end  of  the  delivery  tube  through  a  two-hole  stopper 
into  40  cc.  of  nearly  saturated  Ba(OH)2  solution  contained  in  a  100 
cc.  flask  supported  in  a  large  beaker  of  cold  water.  Boil  in  a  small 
flask  for  about  a  minute  a  mixture  of  25  cc.  water  and  10  cc.  85% 
H3PO4  (to  expel  any  C02  present  in  it).  Pour  this  mixture  into  the 
distilling  flask  with  the  aid  of  a  small  funnel  connected  with  the 
safety-tube.  Heat  the  mixture  to  boiling,  distil  till  about  10  cc. 
have  passed  over,  and  then  remove  the  distillate.  (White  precipitate, 
presence  of  CARBONATE,  SULFITE,  THIOSULFATE,  SULFIDE,  or  SULFUR.) 


p.  Ill 


DETECTION  OF  ACIDIC  CONSTITUENTS. 


103 


Cool  the  distillate  and  make  it  slightly  acid  with  HAc.  (Com- 
plete or  partial  solution  of  the  precipitate,  presence  of  CARBONATE; 
residue  (of  S  or  BaSOs),  presence  in  the  substance  of  free  SULFUR, 
SULFIDE,  SULFITE,  or  THiosuLFATE.)  If  there  is  a  residue,  treat  one- 
half  of  the  mixture  immediately  by  P.  112,  and  separate  portions  of 
the  remainder  by  P.  113,  114,  and  115.  If  there  is  no  residue,  treat 
separate  portions  of  the  whole  distillate  by  P.  113,  114,  and  115. 

Introduce  the  end  of  the  delivery  tube  of  the  distilling  flask  into 
another  receiving  flask  containing  35  cc.  water.  Continue  the  distil- 
lation until  the  liquid  becomes  sirupy,  boils  more  quietly,  and  begins 
to  giv  off  fine  white  fumes.  Treat  this  distillate  as  directed  in  P.  116. 
To  the  contents  of  the  distilling  flask,  while  still  warm,  add  5-10  g. 
of  copper  filings  or  turnings.  Distil  for  3-5  minutes  longer,  col- 
lecting the  distillate  in  15  cc.  of  water.  Note  the  odor  of  the  distillate, 
and  treat  it  by  P.  119. 

Notes. — 1.  It  is  necessary  to  use  a  hard- glass  flask,  since  one  of  ordinary 
glass  is  quickly  destroyed  by  the  action  of  hot,  concentrated  HsPO^  The  boil- 
ing is  sometimes  violent,  especially  when  much  insoluble  material  is  present. 
The  addition  of  the  glass  beads  serves  to  reduce  the  bumping;  and  placing 
the  flask  in  an  inclined  position  prevents  material  from  being  thrown  over 
into  the  distillate,  which  would  lead  to  error  in  the  subsequent  tests.  In  any 
case  in  which  it  seems  possible  that  some  of  the  boiling  liquid  has  been  thrown 
over  into  the  distillate,  a  small  portion  of  the  latter  should  be  tested  for  phos- 
phate by  acidifying  it  with  HNOa  and  adding  an  equal  volume  of  (NEU^MoOi 
solution  (see  P.  104). 

2.  Phosphoric  acid,  which  is  ionized  into  H+  and  H2PO,r  to  a  moderate 
extent  (about  40%  in  0.1  normal  solution),  displaces  almost  completely  from 
their  salts  (unless  these  are  very  difficultly  soluble)  the  much  less  ionized  acids, 


104  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  Ill 

H2C03,  HN02,  H2S,  HC10,  HCN,  HF,  and  HsBO3,  and  also  to  a  large  extent 
the  moderately  ionized  H2SO3.  Since  all  these  acids,  except  HF  and  HaBOs, 
volatilize  readily  out  of  aqueous  solution,  they  pass  over  almost  or  quite  com- 
pletely into  the  first  distillate,  HC1O  in  the  presence  of  chloride  giving  C12. 
The  largely  ionized  acids,  HC1,  HBr,  HI,  HSCN,  HN03,  HC1O3,  H3Fe(CN)«, 
and  H4Fe(CN)e,  are  not  found  in  any  considerable  proportion  in  the  first  dis- 
tillate, since  the  unionized  acid  is  formed  in  much  smaller  proportion,  and 
since  in  addition  it  is  much  less  volatil.  Of  these  the  first  five  pass  over 
unchanged  and  almost  completely  into  the  second  distillate;  for  after  the 
HaPOi  has  become  fairly  concentrated,  the  acids  are  displaced  to  a  greater 
extent  and  volatilize  more  readily  in  consequence  of  the  higher  temperature 
at  which  the  mixture  boils  and  the  smaller  proportion  of  water  it  contains. 
From  the  stronger  HgPC^  solution  HF  also  passes  over  in  large  quantity;  but  this 
is  not  true  of  HaBOs  and  H2SO4,  which  volatilize  only  in  insignificant  amounts 
even  when  the  acid  has  become  nearly  anhydrous.  The  three  acids,  HClOa, 
H4Fe(CN)e,  and  HsFe(CN)6,  are  not  volatil  as  such,  but  are  decomposed  by 
the  H3PO4  after  it  becomes  fairly  concentrated — HC1O3  with  formation  of  Clj 
and  OE,  H3Fe(CN)6  and  H4Fe(CN)6  with  formation  of  free  HCN.  In  regard 
to  the  acids  that  may  be  present  in  the  two  distillates,  see  also  Table  XVI. 

3.  The  barium  salts  of  all  the  acids  passing  into  the  first  distillate,  except 
the  carbonate  and  sulfite,  remain  in  solution.    Phosphoric  acid,  if  thrown  over 
mechanically,  would,  however,  also  giv  a  precipitate.    Sulfur,  when  present  in 
the  free  state  or  when  liberated  from  a  polysulfide  or  thiosulfate,  volatilizes 
with  the  steam,  and  givs  a  turbid  appearance  to  the  water  condensed  in  the 
delivery  tube  and  to  the  barium  hydroxide  solution,  by  which  it  is  little  acted 
on  in  the  cold.     Chlorin  is  converted  by  the  barium  hydroxide  into  barium 
chloride  and  hypochlorite;   bromin,  into  bromide  and  hypobromite,  and  into 
bromide  and  bromate;  and  iodin,  mainly  into  iodate  and  iodide. 

4.  On  acidifying  the  first  distillate  slightly  with  HAc,  BaCOs  dissolves,  but 
BaS03  does  not.    This  difference  in  behavior  is  due  to  the  fact  that  hydrocar- 
bonate-ion    (HCO3~)    is    much    less  ionized    than  hydrosulfite-ion    (HSOa~). 
Sulfur,  if  present,  also  remains  undissolved.    The  addition  of  HAc  causes  the 
liberation  almost  at  once  of  chlorin,  bromin,  or  iodin  from  a  mixture  of  hypo- 
chlorite and  chloride,  hypobromite  and  bromide,  or  iodate  and  iodide;    but 
bromin  is  set  free  somewhat  more  slowly  from  a  mixture  of  bromate  and  bromide. 

5.  A  small  precipitate  obtained  in  this  procedure  (or  in  the  following  one) 
does  not  prove  the  presence  of  carbonate  in  the  mixture  unless  the  prescribed 
precautions  are  carefully  observed — namely,  the  boiling  of  the  original  HsPO4 
solution,  and  avoiding  the  exposure  to  the  air  of  the  various  solutions,  especially 
that  of  the  BaO2H2.    Even  with  these  precautions,  however,  it  is  seldom  pos- 
sible to  prevent  the  absorption  of  enough  CO2  to  produce  a  slight  turbidity. 

6.  Upon  boiling  the  H3PO4  with  the  copper,  H2SO4,  if  present,  is  reduced  to 
H2SO?;   and  this  passes  over  into  the  distillate  in  the  form  of  SO2  gas.    Less 
than  1  mg.  SO4  can  be  detected  by  this  process  of  distillation.     The  copper 
should  be  finely  divided  and  should  be  added  while  the  liquid  is  still  warm, 
since  on  cooling  it  solidifies  to  a  glassy  mass,  which  consists  of  pyrophosphoric 
acid  (H^O?).     The  heating  should  be  continued  for  5-10  minutes;    but,  if 
much  more  prolonged,  the  contents  of  the  flask  change  to  a  solid  mass,  owing 
to  conversion  of  the  pyro  to  metaphosphoric  acid  (HP03),  which  can  after- 
wards be  removed  only  with  much  difficulty. 


P.  112  DETECTION  OF  ACIDIC  CONSTITUENTS.  105 

TABLE  XVII. — ANALYSIS  OP  THE  FIRST  DISTILLATE. 

FIRST  DISTILLATE. — Precipitate:  BaCO3,  BaSO3,  S. 

Solution :  Ba(ClO)2,  Ba(BrO)2,  Ba(IO3)2,  (with  halides) ;  Ba(NO2)2,  BaS,  Ba(CN)2. 
To  the  whole  mixture  add  HAc  (P.  111). 

Precipitate:  BaSO3,  S.     Solution:  H2CO3,  C12,  Br2, 12,  HNCfe,  H2S,  HCN. 
Treat  portions  of  the  unfiltered  mixture  as  follows: 


Add  HCl  and  filter  (P.  112}. 

Add  CHCk 

Filter,  add 

Heat  with 

(P  118) 

Cd(NO  ) 

NaOH  and 

Residue:  S. 

Solution. 

Add  Br2. 

(P.  114). 

FeSOt 

(Shows 

Purple  color  :  I2.  f 

(P.  115). 

SULFIDE 
Or  THIO- 
SULFATE.) 

Precipitate: 
BaSO4.* 

Boil  the  mia 
the  dist 
Ba(C 

Solution: 
H2CO3. 

iture,  collect 
'Hate  in 
>#)2. 

Orange  color:  Br2.J 

Precipitate: 
CdS. 
(Shows 

SULFIDE.) 

Solution: 
Na4Fe(CN)6. 
Precipitate: 
Fe(OH)2-3. 
Add  HCL 

//  the  CHCk  is 
colorless,  add  KI. 
Purple  color:  I2. 
(Shows  NITRITE, 

HYPO  CHLORITE, 

Or   CHLORATE.) 

Blue  pre- 

Precipitate: BaC03. 

cipitate: 

(ShoWS  CARBONATE.) 

Fe4(Fe- 

(CN),),. 

(Shows 

CYANIDE.) 

*  ShoWS  BULFITE  Or  THI08UI.PATB.       fShowS  IODIDE.       J  Shows  BKOMATE  or  BROMIDE. 

Procedure  112. — Detection  of  Carbonate  and  Sulfur-containing  Con- 
stituents.— To  one-half  of  the  first  distillate  (P.  Ill),  if  there  was  a 
residue  on  adding  HAc,  add  1-2  cc.  HCl.  (Residue,  presence  of 
free  SULFUR,  SULFIDE,  or  THIOSULFATE.)  Filter,  and  add  to  the 
filtrate  saturated  bromin  solution  till  the  liquid  becomes  slightly 
yellow.  (White  precipitate,  presence  of  SULFITE  or  THIOSULFATE.) 
Transfer  the  mixture  to  a  distilling  apparatus  such  as  is  used  in  P.  Ill, 
first  filtering  out  the  precipitate  if  it  is  large,  distil  for  a  minute  or 
two,  collecting  the  vapors  in  20  cc.  saturated  Ba(OH)2  solution. 
(White  precipitate,  presence  of  CARBONATE.)  Acidify  slightly  with 
HAc.  (Solution  of  the  precipitate,  presence  of  CARBONATE.) 

Notes. — 1.  See  P.  Ill,  Notes  3-5.  Since  H2SOs  slowly  oxidizes  to  H2SC>4 
in  the  air,  the  solution  should  be  treated  with  HCl  at  once.  If  any  H2SO4 
has  been  formed  in  this  way,  it  will  be  precipitated  as  BaSO4  before  the  addi- 
tion of  Br2.  Care  must  be  taken  to  add  enough  Br2  to  complete  the  oxidation, 
since  otherwise  in  the  subsequent  distillation  SO2  will  distil  over  and  might  be 
mistaken  for  carbonate. 


106  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  112 

2.  If  there  is  a  large  precipitate  of  BaSO4,  it  is  filtered  out,  since  otherwise 
it  is  difficult  to  avoid  violent  bumping  during  the  distillation.     Exposure  to 
the  air,  and  especially  to  the  breath,  should,  however,  be  avoided  so  far  as 
possible,  so  that  CC>2  may  not  be  absorbed  from  it. 

3.  A  residue  of  sulfur  may  arise  from  the  presence  in  the  substance  of  free 
sulfur,  of  a  persulfide,  of  an  ordinary  sulfide  together  with  some  oxidizing 
substance,  or  of  a  thiosulfate. 

Procedure  113. — Detection  of  Nitrite  and  Halogen-liberating  Constit- 
uents.— To  one-fourth  of  what  remains  of  the  first  distillate  (P.  Ill), 
add  1-2  cc.  HAc  and  2-3  cc.  of  chloroform,  and  shake  vigorously. 
(Purple  coloration  of  the  chloroform,  presence  in  the  distillate  of 
free  IODIN;  yellow  or  orange  coloration,  of  free  BROMIN.) 

If  there  is  no  coloration,  add  8-10  drops  of  0.1-normal  KI  solu- 
tion, and  shake  the  mixture.  (Purple  color,  presence  in  the  distillate 
of  CHLORIN,  or  of  NITROUS  ACID;  no  color,  absence  of  NITRITE  and 
HYPOCHLORITE  in  the  substance.) 

If  there  is  a  coloration  either  before  or  after  the  addition  of  KI, 
test  fresh  samples  of  the  original  substance  for  NITRITE,  HYPOCHLO- 
RITE, CHLORATE,  and  BROMATE  by  P.  125,  126,  and  127. 

Notes. — 1.     For  the  reactions  between  the  halogens  and  barium  hydroxide 
and  their  re-formation  on  acidifying  with  HAc,  see  P.  Ill,  Notes  3  and  4. 

2.  The  characteristic  purple  color  given  to  chloroform  is  so  delicate  a  test 
that  even  0.05  mg.  of  iodin  in  the  solution  tested  can  be  detected  by  this  pro- 
cedure.   Bromin  may  be  detected,  but  only  in  the  absence  of  iodin,  by  the 
orange  or  yellow  color  of  the  chloroform  layer  when  not  less  than  0.5  mg.  of 
bromin  is  present  in  the  solution  tested.    Chlorin  gives  no  decided  color  to 
the  chloroform,  but  causes  liberation  of  iodin  on  the  addition  of  KI. 

3.  The  free  halogens  distribute  themselves  between  the  chloroform  and 
water  layers.    In  the  case  of  pure  bromin  or  iodin  the  ratio  of  the  concentration 
in  the  chloroform  to  that  in  the  water  layer  is  very  large;  and  this  ratio  is 
almost  independent  of  the  concentration,  in  accordance  with  the  so-called  dis- 
tribution law,  which  requires  that  the  ratio  of  the  concentrations  of  a  given 
molecular  species,  such  as  Br2  or  1%,  in  any  two  non-miscible  solvents  be  con- 
stant.   When  an  iodide,  like  KI,  is  also  present,  as  it  is  in  the  test  for  free 
chlorin  and  nitrous  acid,  the  proportion  of  iodin  extracted  by  the  chloroform 
is  greatly  reduced,  since  the  iodin  in  the  aqueous  layer  is  largely  combined 
with  the  iodide  in  the  form  of  the  triiodide  (KIs);   but  it  is  still  sufficient 
to  make  the  color-test  a  very  delicate  one,  provided  only  a  few  drops  of  the 
KI  solution  have  been  added. 

4.  For  extracting  the  halogens  from  aqueous  solutions  carbon  tetrachloride 
or  carbon  bisulfide  may  be  used  instead  of  chloroform;  but  carbon  bisulfide  has 
the  disadvantage  of  being  highly  inflammable. 

5.  If  the  tests  described  in  both  paragraphs  of  this  procedure  yield  negativ 
results,  it  shows  the  absence  in  the  substance  of  nitrite  and  hypochlorite,  but 
not  of  the  halides  nor  of  chlorate  and  bromate,  since  these  constituents  may  not 


P.  118  DETECTION  OF  ACIDIC  CONSTITUENTS.  107 


be  decomposed  or  volatilized  till  the  HaPC^  becomes  concentrated  in  the  later 
part  of  the  distillation. 

6.  If  the  chloroform  assumes  a  purple  color  when  it  is  first  added  to  the  dis- 
tillate, it  shows  the  presence  in  the  substance  of  free  iodin,  of  iodate,  or  of  iodide 
(from  which  iodin  has  been  liberated  in  the  distillation  by  the  action  of  the  air 
or  by  some  oxidizing  compound  present  in  the  substance).    If  the  chloroform 
assumes  an  orange  color,  it  shows  the  presence  in  the  substance  of  a  bromate, 
or  of  a  bromide  together  with  some  oxidizing  compound.     In  these  cases  the 
presence  or  absence  of  other  halogen-containing  constituents  and  of  nitrite  has 
to  be  determined  by  the  special  tests  described  in  P.  125-127. 

7.  If  the  chloroform  becomes  colored  only  after  the  KI  is  added,  it  shows  the 
presence  in  the  substance  of  nitrite,  hypochlorite,  or  chlorate  (or  possibly  only 
of  chloride  in  rare  cases  where  a  powerful  oxidizing  substance,  such  as  MnO2, 
KMnO4,  or  K2Cr207,  is  also  present).    Which  one  of  these  constituents  givs  rise 
to  the  color  has  to  be  determined  by  the  special  tests. 

8.  Nitrous  acid  liberates  iodin  from  KI  owing  to  its  reduction  to  nitric 
oxide.     A  peculiarity  of  this  reaction  is  that  the  nitric  oxide  which  is  formed 
by  it  is  rapidly  reoxidized  by  the  oxygen  of  the  air  to  nitrous  acid,  which  then 
reacts  with  the  iodide,  so  that  a  continuous  liberation  of  iodin  results.    Thus 
the  nitrous  acid  acts  as  a  catalyzer  of  the  reaction  between  oxygen  and  HI. 
This  progressiv  liberation  of  iodin  is  highly  characteristic  of  nitrous  acid,  but 
renders  it  difficult  to  estimate  the  amount  of  it  present. 

Procedure  114.  —  Detection  of  Sulfide.  —  To  one-half  of  what  still 
remains  of  the  first  distillate  (P.  Ill),  add  2-3  cc.  Cd(N03)2  solution. 
(Yellow  precipitate,  presence  of  SULFIDE.) 

Note.  —  A  negativ  test  in  the  first  distillate  does  not  prove  the  absence 
of  sulfide  in  the  original  substance,  unless  the  latter  has  dissolved  completely 
in  the  dilute  HsPC^;  for  some  difficultly  soluble  sulfides,  such  as  CuS,  are 
decomposed  only  when  the  HjPOi  becomes  concentrated,  as  it  does  in  the 
latter  part  of  the  distillation.  It  is  therefore  directed  in  P.  116  to  test  also  the 
second  distillate  for  sulfide. 

Procedure  115.  —  Detection  of  Cyanide.  —  Place  what  remains  of  the 

first  distillate  in  a  casserole;    add  1  cc.  NaOH  solution  and  0.5  cc. 

FeSOi  solution;   and  boil  for  one  minute.     To  the  hot  mixture  add 

HC1,  a  few  drops  at  a  time,  until,  on  shaking,  the  dark  colored  pre- 

cipitate of  ferrous  and  ferric  hydroxides  is  dissolved.    Cool  the  mix- 

ture.   If  a  precipitate  is  not  plainly  visible,  filter,  and  v.  ,sh  out  the 

filter-paper  once  with  water.    (Blue  precipitate,  presence  of  CYANIDE.) 

Notes.  —  1.     This  test  is  based  upon  the  formation  of  sodium  ferrocyanide  by 

the  action  of  the  sodium  cyanide  on  the  ferrous  hydroxide  and  upon  the  reaction 

between  this  ferrocyanide  and  the  ferric  salt  which  has  been  produced  by  the 

oxygen  of  the  air.     As  a  result  of  these  two  reactions,  ferric  ferrocyanide 

(Prussian  blue)  is  formed,  which  is  difficultly  soluble  in  dilute  hydrochloric 

acid. 

2.  A  small  precipitate  is  not  readily  detected  in  the  hot  reddish-yellow 
solution,  but  is  more  easily  seen  in  the  cold  light-colored  solution,  especially 


108  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  115 

after  standing,  or  when  collected  on  a  filter.  If  the  precipitate  on  the  filter 
is  not  dark  blue,  it  should  be  washed  with  a  little  hot,  dilute  hydrochloric  acid. 
With  these  precautions,  0.2  mg  CN  in  the  solution  tested  can  be  detected. 

3.  Cyanides  may  be  present  in  the  original  substance  in  the  form  either 
of  simple  or  of  complex  cyanides.     The  latter  are  characterized  by  complex 
anions   (such  as  Ag(CN)?~  and  Fe(CN)6~").     These  differ    very  greatly  in 
their  stability  towards  decomposing  agents,  the  difference  depending  on  the 
extent  to  which  they  are  dissociated  into  the  simple  ions  (Ag+  and  CN~  or 
Fe++  and  CN~).     Ferrocyanides,    ferri cyanides,    and    cobalticyanides  are  so 
slightly  dissociated  in  this  way  that  scarcely  any  HCN  is  produced  when  dilute 
HC1,  HNOs,  or  H2SO4  is  added  to  their  cold  solutions;  but  almost  all  the  other 
complex  cyanides  (such  as  KAg(CN)2  or  K2Ni(CN)4)  are  readily  decomposed  by 
these  acids. — In  the  distillation  with  HsPO^  not  only  the  simple  cyanides,  but 
also  nearly  all  the  complex  cyanides  are  decomposed  during  the  first  part  of 
the  distillation;   but  a  few  very  stable  substances  (such  as  Prussian  blue)  are 
completely  decomposed  only  in  the  second  part  of  the  distillation. 

4.  The  following  procedure  enables  2  mg.  cyanide  to  be  detected  in  the 
presence  of  ferro  or  f erricyanide :    Place  in  a  20  cc.  distilling  flask  provided 
with  a  thistle-tube  0.5-1  g.  of  the  original  substance,  2  g.  powdered  CaCOs, 
and  10  cc.  water.      Add  very  gradually  through  the  thistle-tube  2  cc.  HC1 
(enough  to  decompose  some,  but  not  all,  of  the  CaCOa).    Allow  the  gas  which 
is  evolved  to  pass  into  a  small  test-tube  containing  1  cc.  NaOH  and  5  cc.  water. 
Finally  heat  the  contents  of  the  flask  almost  to  boiling.    Test  the  NaOH  solu- 
tion for  cyanide  by  P.  115. — This  separation  depends  upon  the  fact  that  HCN 
is  displaced  by  ^COs  from  simple  cyanides  and  from  the  relativly  unstable 
complex  cyanides,  such  as  Ag(CN)2~  or  Ni(CN)4=,    but    not  from  ferro  or 
ferricyanides. 

5.  Ferrocyanide  and  ferricyanide  may  be  detected  and  distinguished  from 
each  other  when  only  one  of  them  is  present,  by  adding  a  ferric  salt  to  one 
portion  of  an  aqueous  or  dilute  acid  solution,  and  by  adding  ferrous  salt  to 
another  portion  of  the  solution.    A  ferric  salt  givs  a  blue  precipitate  of  ferric 
ferrocyanide  with  ferrocyanide,  but  no  precipitate  with  a  ferricyanide.     A 
ferrous  salt  givs  the  same  blue  precipitate  (of  ferric  ferrocyanide)  with  a  ferri- 
cyanide;   but  it  also  givs  with  a  ferrocyanide  a  precipitate  (of  ferrous  ferro- 
cyanide), which  is  white  if  no  ferric  salt  is  present,  but  which  rapidly  turns 
blue  in  contact  with  the  air. 

6.  Ferrocyanide  and  ferricyanide  may  be  detected  in  the  presence  of  each 
other  by  proceeding  as  follows:  Add  to  an  aqueous  or  dilute  HNOs  solution  of  the 
substance  AgNOs  and  then  a  moderate  excess  of  NEUOH.      (White  precipitate 
insoluble  in  NELiOH,  presence  of  FERROCYANIDE.     Orange  to  red  precipitate 
readily  soluble  in  NKLiOH,  presence  of  FERRICYANIDE.)     Filter  out  and  wash 
the  precipitate,  and  pour  over  it  a  little  FeCls  solution.     (Blue  coloration, 
presence  of  FERROCYANIDE.)    Acidify  the  ammoniacal  filtrate  with  HAc,  filter 
out  and  wash  the  precipitate,  and  pour  through  the  filter  containing  it  a  little 
FeSO4  solution.     (Orange-red  precipitate,  which  is  turned  blue  by  the  FeSO4, 
presence  of  FERRICYANIDE).    This  procelure  enables  0.2  mg.  Fe(CN)e  as  either 
ferro  or  ferricyanide  to  be  detected  when  present  alone;   but  the  test  for  ferri- 
cyanide is  much  less  delicate  in  the  presence  of  much  ferrocyanide. 


P.  116  DETECTION  OF  ACIDIC  CONSTITUENTS.  109 

TABLE  XVIII. — ANALYSIS  or  THE  SECOND  AND  THIRD  DISTILLATES. 


SECOND  DISTILLATE:  H2S,  HCN,  HSCN,  HC1,  HBr,  HI,  C12,  Br2,  I2. 

THIRD  DIS- 
TILLATE: 
H2SOs. 
Add  HCl, 
BaCk,  and 
Brz(P.119). 

To  a  portion 
add  AgN03 
(P.  116). 

//  AgNOs  givs  a  precipitate,  treat  portions  as  follows: 

Add 
Cd(NOs)2 
(P.  W). 

Add  FeCk 
and  HCl 
(P.  117). 

Add  CHCk  (P.  118). 

Precipitate: 
AgCl,  AgBr, 
Agl,  AgSCN, 
Ag2S, 

Ag2(CN)2. 

CHCla  layer: 
I2,  Br2,  C12.* 
//  colorless, 
add  KL 

Water  layer: 
HCl,  HBr,  HI. 

Test  for  sepa- 
rate HALIDES 
by  P.  106. 

Precipitate: 
BaSO4. 

(Shows 

SULFATE.) 

Precipitate: 
CdS. 
(Shows  SUL- 

FIDE.) 

Red  color: 

Fe(SCN)3. 
(Shows  THIO- 

CYANATE.) 

Purple  color: 
I2.     (Shows 

CHLORATE 
Or  CHLORIDE.) 

*  Purple  coloration  shows  IODIDE,  orange  coloration,  BROMIDE  or  BROMATE. 

Procedure  116. — Detection  of  Constituents  Precipitable  by  Silver 
Nitrate. — To  one-sixth  of  the  second  distillate  (P.  Ill)  add  1  cc. 
HN03  and  1  cc.  AgNOs  solution.  (White  precipitate,  presence  of 
CHLORIDE,  CYANIDE,  or  THiocYANATE;  yellowish  precipitate,  presence 
of  BROMIDE  or  IODIDE;  black  precipitate,  presence  of  SULFIDE.) 

If  there  is  a  precipitate,  test  one-sixth  of  the  second  distillate  for 
sulfide  by  P.  114,  another  sixth  for  thiocyanate  by  P.  117,  and  the 
remainder  for  free  halogen  and  halides  by  P.  118  followed  by  P.  106. 
If  there  is  no  precipitate,  reject  the  whole  second  distillate. 

Notes. — 1.  As  to  the  solubilities  of  the  various  silver  salts,  see  Note  2,  P.  103. 
2.  It  is  not  necessary  to  test  for  cyanide  in  this  distillate;  for  even  the 
insoluble  ferro  and  ferricyanides  are  decomposed  partly,  tho  not  necessarily 
completely,  in  the  first  part  of  the  distillation.  It  is,  however,  advisable  to 
test  for  sulfide  unless  the  substance  dissolved  completely  in  the  hot  dilute 
HaPO,!,  or  unless  the  AgNOa  precipitate  is  pure  white;  for  some  insoluble 
sulfides  begin  to  decompose  only  when  the  HsPOi  becomes  concentrated. 

Procedure  117. — Detection  of  Thiocyanate. — Dilute  a  sixth  of  the 
second  distillate  (P.  Ill)  to  5-10  cc.,  add  2-3  drops  of  FeCl3  solu- 
tion and  2-3  drops  of  HCl.     (Red  color,  presence  of  THIOCYANATE.) 
Notes. — 1.     The  red  coloration  arises  from  the  formation  by  metathesis  of 
Fe(SCN)s,  a  substance  whose  degree  of  ionization  is  relativly  small.    The  HCl 
is  added  to  reduce  the  hydrolysis  of  the  FeCls  and  diminish  the  color  imparted 
by  it  to  the  solution. — A  distinct  reddish-yellow  coloration  is  produced  by 
0.1  mg.  SCN.    A  deep  red  color  is  obtained  when  1  mg.  or  more  is  present. 

2.  Since  in  the  distillation  with  HsPC^,  thiocyanates  are  destroyed  by 
certain  oxidizing  agents,  such  as  nitrates,  it  is  sometimes  advisable  to  apply 
this  test  also  to  a  solution  of  the  original  substance. 


110  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  118 

Procedure  118. — Detection  of  Chlorate,  Bromate,  and  the  Halides. — If 
AgNO3  gave  a  precipitate  in  P.  116,  to  the  remainder  of  the  second 
distillate  in  a  separating  funnel  add  2-3  cc.  chloroform  and  shake. 
(Purple  color,  presence  of  IODIDE;  orange  or  yellow  color,  presence  of 

BROMIDE  Or  BROMATE.) 

If  the  chloroform  is  colorless,  separate  it  from  the  aqueous  layer; 
and  add  to  the  chloroform  layer  a  few  drops  of  KI  solution.  (Purple 
color,  probable  presence  of  CHLORATE.)  If  there  is  no  color,  treat  the 
aqueous  layer  left  in  the  separating  funnel  by  P.  106. 

If  the  chloroform  becomes  colored  either  before  or  after  the  addi- 
tion of  KI,  add  to  the  mixture  in  the   separating  funnel  enough 
H2S03  solution  to  reduce  the  halogen,  draw  off  the  chloroform  layer 
if  it  is  still  in  the  funnel,  and  treat  the  aqueous  layer  by  P.  106. 
Notes. — 1.    As  to  these  tests  see  the  notes  to  P.  113. 

2.  In  the  second  part  of  the  distillation  chlorate  or  bromate,  whether 
present  alone  or  with  a  halide,  is  rapidly  decomposed  with  evolution  of  Ch  or 
Br2.  Therefore,  if  free  halogen  is  found  present  neither  in  the  first  nor  second 
distillate,  it  shows  the  absence  of  chlorate  or  bromate.  The  presence  of  Cl2  or 
Br£  in  the  second  distillate  does  not,  however,  necessarily  indicate  chlorate  or 
bromate;  for  these  halogens  may  be  produced  from  the  corresponding  halides 
by  the  action  of  some  oxidizing  substance.  The  special  test  for  chlorate  or 
bromate  described  in  P.  127  should  therefore  be  tried  when,  and  only  when, 
free  halogen  is  found  in  either  the  first  or  second  distillate. 

Procedure  119. — Detection  of  Sulfate. — To  the  third  distillate  ob- 
tained upon  heating  with  copper  (P.  Ill),  add  1-2  cc.  HC1,  3-5  cc. 
BaCl2  solution,  and  saturated  Br2  solution  till  the  liquid  becomes 
yellow.  (White  precipitate,  presence  of  SULFATE.) 

Notes. — 1.  By  the  action  of  copper  in  the  presence  of  concentrated  H^PC^  on 
sulfates  (even  on  the  very  difficultly  soluble  BaSO4)  SO2  is  formed.  This  is 
oxidized  by  the  Br2  to  H2S04,  which  then  precipitates  as  BaS(>4.  In  this  way 
1  nig.  SC>4  may  be  detected.  Even  when  this  small  amount  is  present  in  the 
substance,  only  a  small  proportion  of  it  passes  into  the  second  distillate. 

2.  Much  HaPOi  also  passes  over  into  the  distillate;  and  the  HCi  is  added 
to  prevent  its  precipitation  as  BaHPC>4.    Too  much  HCI  must  not  be  added 
since  BaS(>4  is  appreciably  soluble  in  it. 

3.  When  a  sulfide  is  present  which  has  not  already  been  decomposed,  sulfur 
and  H2S,  or  sulfur  and  SC>2,  may  pass  into  the  third  distillate,  after  the  acid 
has  become  concentrated.     The  H2S  may  be  tested  for  in  a  portion  of  the 
distillate  by  P.  114.     Owing  to  the  possible  formation  of  SOz  by  the  oxidation 
of  sulfur  at  the  high  temperature  attained  at  the  end  of  the  distillation,  the 
test  for  sulfate  is  unreliable  if  sulfur,  undecomposed  sulfide,  or  S02  gas  is 
present  in  the  distilling  flask  at  the  end  of  the  second  part  of  the  distillation. 
In  that  case  a  fresh  0.3  g.  sample  of  the  substance  should  be  heated  with  10  cc. 
2-nonnal  HCI,  the  mixture  filtered,   and   the  filtrate   tested  for  sulfate  by 
adding  an  equal  volume  of  BaCl2  solution. 


P.  121 


DETECTION  OF  ACIDIC  CONSTITUENTS. 


Ill 


SUBSTANCES    DECOMPOSED  BY  COLD  DILUTE  ACIDS  OR  BY  HOT    CONCEN- 
TRATED ACIDS!    SUPPLEMENTARY  PROCEDURES. 


TABLE  XIX. — SUPPLEMENTARY  PROCEDURES  FOR  DETECTING  THE  ACIDIC 

CONSTITUENTS. 


Treat  samples  of  the  original  substance  as  follows: 


Distil  with 
CHzOH  and 
#2/S04 
(P.  121). 

Heat  with 
SiOz  and 
KHSO* 
(P.  122). 

Boil  with 
HN03,  add 
(NHJtMoOt 
(P.  123). 

Distil  with 
HzSOi  and 
FeSO* 
(P.  124). 

Dissolve  in 
water,  intro- 
duce into  an 
inverted  tube 
filled  with 
HCl  solution 
of  urea 
(P.  126). 

Heat  with 
water,  HAc, 
and  PbAcz 
(P.  126). 

Dissolve  in 
HN03, 
add  AgNOs 
(P.  127). 

Distillate: 
B(OCH»)3. 
Collect  in 
CHSOH 
and  HCl, 
and  add 
turmeric. 

Gases 
evolved: 
SiF4  and 
H2O. 
Deposit  on 
cold  part 
of  tube: 
Si03H2. 
(Shows 

FLUORIDE.) 

Yellow 
precipitate  : 
(NH4)3P04. 
12MoO3. 
(Shows 

PHOSPHATE.) 

Distillate: 
HNO2. 
Add  KI  and 
CPICk. 

Dark  brown 
precipitate  : 
Pb02. 

(Shows 

HYPOCHLO- 
RITE.) 

Precipitate: 
AgCl,  etc. 
Reject. 

Gas:  N2. 
(Shows 

NITRITE.) 

Filtrate: 
AgC103. 
Add  H2SOZ 

Purple  color: 
I2. 
(Shows 

NITRATE  Or 
NITRITE.) 

Precipitate: 
AgCl. 
(Shows 

CHLORATE.) 

Orange  color. 
(Shows 

BORATE.) 

Procedure  121. — Detection  of  Borate. — Place  1-2  g.  of  the  finely 
powdered  substance  in  the  distilling  apparatus  used  in  P.  Ill,  and 
add  10  cc.  of  methyl  alcohol  (CH3OH)  and  two  or  three  glass  beads. 
Pour  in  carefully  3  cc.  96%  H2SO4,  and  distil  off  most  of  the  alcohol, 
collecting  it  in  a  mixture  of  5  cc.  CH3OH  and  3  cc.  12-normal  HCl. 
Cool  the  distillate,  and  add  to  it  five  drops  of  a  saturated  solution  of 
turmeric  in  ethyl  alcohol.  (Red  or  orange  color,  presence  of  BORATE.) 

Note. — Methyl  alcohol  reacts  with  boric  acid  to  form  methyl  borate  B(OCH3)a 
which  is  a  readily  volatil  liquid.  The  color  given  by  turmeric  to  a  solution  of 
boric  acid  in  methyl  alcohol  and  strong  hydrochloric  acid  is  so  intense  that  the 
test  is  very  delicate  if  the  proportions  given  are  reproduced.  The  presence  of 
1  mg.  BO2  in  the  substance  distilled  may  readily  be  detected.  To  estimate 
roughly  the  quantity  present,  the  color  may  be  compared  with  that  given  by 
adding  the  turmeric  solution  to  known  quantities  of  borate  dissolved  in  a 
mixture  of  3  cc.  12-normal  HCl  and  15  cc.  CHsOH. 

Procedure  122. — Detection  of  Fluoride. — Mix  0.2  g.  of  the  dry, 
finely  powdered  substance  with  twice  its  weight  of  powdered  KHSOi 
and  with  10-20  mg.  dry,  finely  powdered  or  precipitated  Si02.  Blow 


112  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  122 

a  thick-walled  bulb  1J-2  cm.  in  diameter  at  the  end  of  a  glass  tube 
of  5-8  mm.  bore.  Place  the  mixture  in  the  bulb  (not  using  more  of 
it  than  will  one-third  fill  the  bulb).  Heat  the  bulb  carefully  until 
the  KHSO4  is  melted,  taking  care  that  the  mixture  does  not  froth 
up  into  the  tube.  Continue  to  heat  the  bulb  and  the  lower  part  of 
the  tube  until  there  is  a  deposit  of  a  solid  substance  or  of  condensed 
acid  3  or  4  cm.  above  the  bulb.  After  it  has  cooled,  cut  off  the  tube 
close  to  the  bulb.  Dip  the  tube  several  times  in  water,  dry  it  in  a 
flame,  and  heat  it  strongly.  (White  deposit  in  the  middle  part  of 
the  tube  and  etched  surface  at  the  lower  end,  presence  of  FLUORIDE.) 

Notes. — 1.     This  test  depends  on  the  following  reactions: 
4HF  +SiO2  =SiF4     +2H2O. 
3SiF4+3H2O  =H2SiOs+2H2SiF6. 

Some  of  the  HF  liberated  by  the  molten  KHSO4  volatilizes  and  takes  the  silica 
required  for  the  first  reaction  from  the  glass,  thus  producing  the  characteristic 
etched  surface  in  the  lower  part  of  the  tube.  The  SiF4  gas  and  the  water- 
vapor  liberated  react  in  the  cooler  part  of  the  tube  according  to  the  second 
equation  (forming  a  white  ring  of  solid  silicic  acid  and  fluosilicic  acid,  H2SiF«). 
The  reaction  is  reversed  at  higher  temperatures,  so  that  the  deposit  may  be 
driven  up  the  tube  by  heating.  This  white  deposit  is  the  most  characteristic 
part  of  the  test  for  fluoride.  A  deposit  of  SOs  and  H^SO*  may  also  form  in  the 
upper  part  of  the  tube,  and  might  be  mistaken  for,  or  interfere  with,  the  test 
for  small  amounts  of  fluoride,  if  the  final  washing  with  water  is  omitted.  This 
procedure  enables  0.5  mg.  F  to  be  easily  detected. 

2.  The  test  fails  with  certain  minerals  which  are  not  decomposed  by  fusion 
with  KHSOi.     Such  cases  are  provided  for  by  the  treatment  described  in  P.  131. 

3.  Fluoride  is  often  tested  for  by  heating  the  solid  substance  in  a  platinum 
crucible  with  H2S(>4  alone  and  detecting  any  HF  evolved  by  its  etching  action 
on  a  watch  glass  coated  with  wax  through  which  markings  have  been  made. 
This  test  has  the  disadvantage  that  when  silica  or  silicate  is  present,  which  is 
very  often  the  case  in  minerals,  it  is  unreliable  owing  to  the  conversion  of  the 
HF  to  SiF4  by  the  reaction  given  in  Note  1. 

Procedure  123. — Detection  of  Phosphate. — To  0.1-0.2  g.  of  the 
finely  powdered  substance  add  about  5  cc.  HNOs.  If  the  substance 
does  not  dissolve,  boil  the  mixture  for  2  or  3  minutes,  and  filter. 
Add  to  the  filtrate  an  equal  volume  of  (NH4)2Mo04  solution,  and 
allow  it  to  stand  5  to  10  minutes.  (Yellow  precipitate,  presence  of 
PHOSPHATE.) 

Note. — See  the  notes  on  P.  104. 

Procedure  124. — Detection  of  Nitrate  or  Nitrite. — Arrange  a  distil- 
lation-apparatus in  the  way  shown  in  the  figure  under  P.  111. 
Place  in  the  distilling  flask  0.3  g.  of  the  solid  substance,  5  cc. 
2-normal  FeSC>4  solution,  and  15  cc.  H2S04;  and  place  in  the  receiv- 
ing flask  a  mixture  of  20  cc.  water  and  1  cc.  NaOH  solution.  Distil 


P.  124  DETECTION  OF  ACIDIC  CONSTITUENTS.  113 

till  only  about  5  cc.  remain  in  the  distilling  flask.  Acidify  the 
distillate  with  H2S04,  add  2-3  cc.  chloroform,  and  shake  (to  make 
sure  that  the  chloroform  remains  colorless).  Then  add  8-10  drops  of 
KI  solution,  and  shake  again.  (Purple  coloration  of  the  chloroform, 
presence  of  NITRATE  or  NITRITE.) 

Notes. — 1.  In  this  procedure  the  nitrate  is  reduced  by  the  FeSC>4  to  nitric 
oxide  (NO),  which  passes  over  as  a  gas  into  the  receiver,  where  it  is  oxidized 
by  the  oxygen  of  the  air  to  HNO2,  which  is  then  absorbed  by  the  NaOH.  When 
the  solution  is  acidified  and  KI  added,  I2  is  liberated  by  the  HNC>2.  By  this 
procedure  0.2  mg.  NOs  can  be  detected. 

2.  The  reaction  is  highly  characteristic  for  nitrates  and  nitrites,  since  other 
oxidizing  substances  (for  example,  chlorin  or  bromin)  which  might  liberate  iodin 
from  potassium  iodide  are  reduced  by  the  FeSCU  to  compounds  which,  even  if 
they  pass  over  into  the  distillate,  have  no  action  on  KI.  The  only  substances 
that  may  interfere  are  iodides  and  thiocyanates ;  if  these  are  present,  they  should 
be  removed  before  the  distillation  by  treating  the  substance  with  10  cc.  H2SO4, 
adding  solid  Ag2SO4,  shaking,  and  filtering. 

Procedure  125. — Detection  of  Nitrite. — To  0.1  g.  of  the  substance 
add  1  cc.  of  water  and  4-5  drops  of  HC1.     Fill  a  7-cm.  test-tube 
with  a  20%  solution  of  urea  in  HC1,  and  invert  it  over  a  small  dish 
containing  more  of  the  same  solution.    Introduce  the  solution  of  the 
substance  into  the  test-tube  by  means  of  a  small  tube  which  has  one 
end  closed  with  a  rubber  nipple  and  the  other  end  drawn  out  and 
bent  so  as  to  form  a  small  U.     Take  care  not  to  introduce  an  air- 
bubble  at  the  same  time.     (Formation  of  gas,  presence  of  NITRITE.) 
Notes. — 1.     The  reaction  between  urea  and  nitrous  acid  is 
CO  (NH2)2  -I-  2HNO2  =  CO2  +2N2  +3H2O. 

The  N2  is  liberated  hi  the  form  of  minute  bubbles  which  collect  at  the  top 
of  the  tube.  When  much  CO2  is  produced,  it  also  separates  as  a  gas;  but  a 
small  quantity  remains  dissolved  hi  the  liquid. 

2.  This  procedure  enables  0.1  mg.  NO2  to  be  detected.     The  amount  of 
nitrite  present  may  be  estimated  by  making  a  comparativ  test  with  a  known 
quantity  of  nitrite. 

3.  The  mixture  is  acidified  before  it  is  introduced  into  the  tube  so  that  any 
carbonate  present  may  be  expelled  from  it. 

4.  The  halogens,  chlorin,  bromin,  and  iodin,  when  dissolved  in  alkali,  de- 
compose urea  with  evolution  of  nitrogen,  but  they  do  not  do  so  when  dissolved 
in  concentrated  HC1.     They  do  not  therefore  interfere  with  the  test  when 
carried  out  as  above  described. 

Procedure  126. — Detection  of  Hypochlorite. — To  0.5  g.  of  the  pow- 
dered substance  add  5  cc.  water,  and  then  HAc,  a  few  drops  at  a  time, 
until  the  solution  is  acid.  Filter  if  there  is  much  residue,  add  2-3  cc. 
PbAca  solution,  heat  the  mixture  to  boiling,  and  let  it  stand  for  ten 
minutes.  (Brown  precipitate,  presence  of  HTPOCHLORITE.) 


114  DETECTION  OF  ACIDIC  CONSTITUENTS.  P.  126 

Notes. — 1.  Hypochlorites  are  commonly  met  with  either  in  alkaline  solution 
or  in  the  form  of  a  powder  (for  example,  in  bleaching  powder).  Since  they 
are  prepared  by  the  action  of  chlorin  on  alkali,  chloride  is  ordinarily  present 
in  nearly  equivalent  amount.  When  the  solid  powder  is  treated  with  water, 
the  hypochlorite  passes  into  solution;  and  from  it  the  unionized  HC1O  is 
liberated  upon  the  addition  of  the  more  largely  ionized  acetic  acid.  Chlorin 
is  also  formed  in  such  quantity  as  will  satisfy  the  equilibrium-conditions  of  the 
reaction  HC1O+C1-+H+=C12+H2O.  When  in  neutralizing  with  HAc  litmus 
paper  is  used,  the  paper  will  soon  be  bleached  if  hypochlorite  is  present;  but 
the  color  at  the  first  instant  or  on  the  edges  of  the  bleached  portion  can  usually 
be  observed. 

2.  This  test  depends  upon  the  oxidation  of  the  lead  salt  to  lead  dioxide 
(PbC^)  by  the  hypochlorite.    The  reaction  takes  place  so  slowly  in  the  cold  that 
not  less  than  10  mg.  CIO  in  5  cc.  solution  can  be  detected  at  room  tempera- 
ture, even  if  the  mixture  be  allowed  to  stand  a  few  minutes.    But  when  the 
mixture  is  heated,  the  limit  of  detectability  is  about  0.5  mg.  in  5  cc.    The  solu- 
tion is  acidified  with  HAc,  since  oxidation  does  not  take  place  in  the  presence 
of  a  strong  acid,  such  as  HNOs. 

3.  Peroxides  in  alkaline  solution  react  instantaneously  with  lead  salts, 
forming  PbO2;   but  this  reaction  does  not  take  place  in  the  presence  of  HAc, 
even  on  boiling.    Therefore  in  the  above  procedure  a  peroxide  will  not  be  mis- 
taken for  a  hypochlorite.     Peroxide  and  hypochlorite,  moreover,  cannot  exist 
together,  since  they  react  very  rapidly  with  formation  of  oxygen. 

4.  This  test  for  hypochlorite  may  be  made  even  more  delicately  in  alkaline 
solution,  provided  peroxides  are  known  to  be  absent.     If  the  solution  is  only 
slightly  alkaline,  a  small  white  precipitate  of  Pb(OH)2  or  PbCOs  is  first  formed; 
but  this  turns  brown  if  hypochlorite  is  present  when  the  mixture  is  heated  and 
allowed  to  stand.     The  delicacy  is  of  course  diminished  by  the  presence  of 
a  large  amount  of  Pb(OH)2  or  PbCOa;  but  1  mg.  CIO  can  be  detected  in  the 
presence  of  even  2  or  3  g.  of  these  substances,  provided  an  excess  of  the  lead 
salt  is  still  present  in  the  solution  and  the  mixture  is  boiled  vigorously,  prefer- 
ably in  a  casserole. 

Procedure  127. — Detection  of  Chlorate  and  Br ornate. — Treat  0.3  g. 
of  the  powdered  substance  in  the  cold  with  30  cc.  water.  (If  hypo- 
chlorite is  present  as  shown  by  P.  126,  reduce  it  to  chloride  by  adding 
NaAs02  solution  till  a  drop  of  the  mixture  placed  on  filter-paper  wet 
with  starch  and  KI  solution  no  longer  produces  a  blue  coloration.) 
Add  20  cc.  HN03  and  10  cc.  AgN03  solution.  Shake  the  mixture, 
and  filter  off  the  precipitate.  To  the  filtrate  add  5  cc.  of  saturated 
862  solution,  and  allow  it  to  stand  five  minutes.  If  there  is  a  pre- 
cipitate, heat  the  mixture  nearly  to  boiling.  (White  precipitate, 
presence  of  CHLOKATE;  yellowish  precipitate,  presence  of  BROMATE.) 
Note.— See  Notes  3  and  4,  P.  103. 


P.  131  DETECTION  OF  ACIDIC  CONSTITUENTS.  115 

SUBSTANCES    NOT    DECOMPOSED    BY    CONCENTRATED    ACIDS. 


Procedure  131. — Detection  of  All  the  Acidic  Constituents. — Treat 
two  0.3  g.  samples  of  the  powdered  substance  by  P.  101;  or  treat 
a  2  g.  sample  of  it  by  the  first  two  paragraphs  of  P.  Ill,  testing 
portions  of  the  distillate  thus  obtained  by  P.  112-115. 

Treat  2  g.  of  the  very  finely  powdered  substance  with  30  cc.  of 
cold  1-normal  HNOs,  filter,  and  wash  the  residue.  Treat  portions 
of  the  solution  by  P.  102-106.  Separate  the  residue  from  the  filter, 
if  possible;  incinerating  the  paper,  if  much  residue  adheres  to  it,  in 
a  spiral  of  platinum  wire.  Transfer  the  residue  (with  the  ash)  to  a 
nickel  crucible,  and  dry  it  by  igniting  it  gently.  Mix  it  in  the 
crucible  with  10-15  g.  dry  Na2C03.  Cover  the  crucible,  and  heat  it, 
preferably  within  a  cylindrical  jacket  of  asbestos-paper,  over  a  power- 
ful burner  for  15-20  minutes.  If  a  perfectly  clear  fusion  does  not 
result,  add  more  Na2C03,  and  heat  again.  Cool,  place  the  crucible 
in  a  casserole,  boil  it  with  water  till  the  fused  mass  is  disintegrated, 
and  filter,  rejecting  the  residue. 

To  three-fourths  of  the  filtrate  add  HNOs  till  it  is  distinctly  acid, 
then  add  5  cc.  more,  filter  if  there  is  a  precipitate  (of  silicic  acid), 
and  test  portions  of  the  solution  for  sulf ate,  fluoride,  sulfide,  halides, 
and  phosphate,  by  P.  102,  103,  and  104. 

Test  the  remainder  of  the  filtrate  for  borate  by  evaporating  it  to 
complete  dryness,  adding  96%  H2S04  drop  by  drop  as  long  as  there 
is  effervescence,  and  treating  the  mixture  by  P.  121. 

If  silicate  is  present,  test  a  fresh  sample  of  the  substance  for 
fluoride  by  P.  122. 

Notes. — 1.    Fusion   with   NaaCOa   metathesizes   nearly   all   insoluble   com- 
pounds in  the  way  described  in  the  Notes  to  P.  7. 

2.  In  minerals  or  metallurgical  products  undecomposed  by  acids,  it  is 
usually  necessary  to  test  only  for  silicate,  chloride,  sulfate,  phosphate,  borate, 
and  fluoride,  since  other  acidic  constituents  are  scarcely  ever  present. 

3.  In  the  process  of  fusion  changes  in  the  state  of  oxidation  may  take  place. 
Thus  sulfate  may  be  wholly  or  partially  reduced  to  sulfide,  and  sulfide  wholly 
or  partially  oxidized  to  sulfate.     These  two  constituents  must  therefore  be 
distinguished  by  tests  made  with  the  unfused  substance,  as  described  in  the 
first  part  of  this  Procedure. 


APPENDICES 


PREPARATION  OF  THE  REAGENTS 


ACIDS. 

Acetic,  6-normal:  Mix  350  cc.  99.5%  acid  with  650  cc.  water. 
Hydrochloric,  12-normal:  Use  the  C.  P.  acid  of  commerce  of  s.  g.  1.19. 
Hydrochloric,  6-normal:   Mix  12-normal  HC1  with  an  equal  volume 

of  water. 

Hydrofluoric,  48  percent. :   Use  the  pure  acid  sold  in  ceresin  bottles. 
Nitric,  16-normal:  Use  the  C.  P.  acid  of  commerce  of  s.  g.  1.42. 
Nitric,  6-normal:   Mix  380  cc.  HNO3  (s.  g.,  1.42)  with  620  cc.  water. 
Perchloric,  2-normal:  Use  the  purest  acid  of  commerce  of  s.  g.  1.12. 
Phosphoric,  85  percent. :  Use  the  C.P.  acid  of  commerce. 
Sulfuric,  96  percent.:  Use  the  C.  P.  acid  of  commerce  of  s.  g.  1.84. 
Sulfuric,  6-normal:    Pour  96%  H2S04  into  five  volumes  of  water. 
Sulfurous,  saturated:    Saturate  water  at  20-25°  with  S02  gas  made 

by  dropping  96%  H2S04  into  hot  NaHS03  solution. 
Tartaric,  10  percent.:  To  100  g.  of  the  solid  add  enough  water  to 

make  1000  cc.  of  solution. 

BASES. 

Ammonium  hydroxide,  15-normal:  Use  the  C.P.  product  of  s.  g.  0.90. 
Ammonium  hydroxide,  6-normal:    Mix  400  cc.  15-normal  NH4OH 

with  600  cc.  water. 
Barium  hydroxide,  saturated:  Heat  60  g.  Ba(OH)2.8H20  with  1000  cc. 

water,  cool  to  15°,  decant  or  filter. 
Sodium  hydroxide,  6-normal:  Add  to  250  g.  NaOH   "purified  by 

alcohol"  enough  water  to  make  the  volume  1000  cc. 

AMMONIUM   SALTS. 

Acetate,  3-normal :  Mix  equal  volumes  of  6-normal  HAcand  6-normal 

NH4OH;  or  dissolve  250  g.  of  the  solid  salt  in  enough  water 

to  make  the  volume  1000  cc. 
Carbonate:    Dissolve  250  g.  freshly  powdered  ammonium  carbonate 

in  1000  cc.  6-normal  NH4OH,  and  filter  if  there  is  a  residue. 
Chloride,  1-normal:   Add  to  54  g.  NH4C1  enough  water  to  make  the 

volume  1000  cc. 
Molybdate:    Dissolve  75  g.  of  the  pure  ammonium  molybdate  of 

commerce  in  500  cc.  water,  pour  the  solution  into  500  cc. 

6-normal  HN03,  and  shake  the  mixture  occasionally  till  the 

precipitate  is  dissolved. 
Monosulfide,  6-normal:  Pass  H2S  gas  into  200  cc.  15-normal  NH4OH 

in  a  bottle  immersed  in  running  water  or  in  iced  water  until 

the  gas  is  no  longer  absorbed;   then  add  200  cc.  15-normal 

NH4OH  and  enough  water  to  make  the  volume  1000  cc. 

117 


118 


APPENDICES 


Oxalate,  0.5  normal:  Dissolve  35  g.  (NH4)2C2O4.H20  in  1000  cc.  water. 

Polysulfide,  6-normal:    Digest  one  liter  of  ammonium  monosulfide 

with  25  g.  flowers  of  sulfur  for  some  hours  and  filter. 


Name  of  SaU 

Bariuin  chloride 
Cadmium  nitrate 
Calcium  chloride 
Calciu^m  sulfate 
Cobalt  nitrate 
Copper  Sulphate 
Ferric  chloride 
Lead  acetate 
Mercuric  chloride 
Potassium  chromate 
Potassium  cyanide 
Potassium  ferricyanide 
Potassium  ferrocyanide 
Potassium  iodide 
Potassium  nitrite 
Potassium  permanganate 
Potassium  thiocyanate 
Silver  nitrate 
Sodium  acetate 
Sodium  arsenite 
Sodium  carbonate 
Sodium  nitrite 
Sodium  phosphate 


OTHER   SALTS. 

Formula  and 
formula-weight 
BaCl2.2H20(244) 
Cd(N03)2.4H2O(308) 
CaCl2.6H2O(219) 
CaS04.2H20(172) 
Co(N03)2.6H20(291) 
CuSO4.5H2O(250) 
FeCl3.6H2O(270) 
Pb(C2H3O2)2.3H20(379) 
HgCl2  (271) 
K2Cr04  (194) 
KCN  (65) 
K3Fe(CN)6  (329) 
K4Fe(CN)6.3H20(422) 
KI  (166) 
KN02  (85) 
KMnO4  (158) 
KSCN  (97) 
AgNOs  (170) 
NaC2H302.3H20(136) 
NaAs02  (130) 
Na2C03  (106) 
NaNO2(69) 
Na2HPO4.12H2O(358) 


Concen-        Grams 


tration 
1 -normal 
1-normal 
1-normal 
saturated 
1%  cobalt 
1-normal 
1-normal 
1-normal 
0.2-normal 
3-normal 
1-normal 
1-normal 
1-normal 
0.1 -normal 
3-normal 
1-per  cent. 
1-normal 
1-normal 
1-normal 
1-normal 
3-normal 
3-normal 
1-normal 


per  liter 
120 
150 
110 
2 

50 
125 

90 
190 

25 
290 

65 
110 
105 

17 
250 

10 
100 
170 
135 
130 
160 
210 
120 


SPECIAL   REAGENTS. 

Bromin,  saturated  solution:  Shake  liquid  bromin  with  water,  leaving 
a  small  excess  of  it  in  contact  with  the  solution. 

Ether  saturated  with  HC1:  pass  dry  HC1  gas,  as  long  as  it  continues 
to  be  absorbed,  into  a  bottle  of  ether  immersed  in  ice-water. 

Ferrous  sulfate,  2-normal:  Dissolve  280  g.  FeS04.7H2O  in  0.6-normal 
H2S04,  and  keep  in  contact  with  iron  nails. 

Hydrogen  peroxide,  3  per  cent. 

Magnesium  ammonium  chloride,  1-normal  in  MgCl2:  Dissolve  100  g. 
MgCl2.6H2O  and  100  g.  NH4C1  in  water,  add  50  cc.  15- 
normal  NH4OH,  and  dilute  to  1000  cc. 

Potassium  mercuric  iodide,  0.5  normal  in  K2HgIi:  Dissolve  115  g. 
HgI2  and  80  g.  KI  in  enough  water  to  make  the  volume  500 
cc. ;  add  500  cc.  6-normal  NaOH ;  and  decant  the  solution  from 
any  precipitate  that  may  form  on  standing.  Keep  this  stock 
solution  in  the  dark. 


APPENDICES. 


119 


Potassium  pyroantimonate :  Add  20  g.  of  the  best  commercial  salt  to 
1000  cc.  boiling  water,  boil  for  a  minute  or  two  till  nearly  all 
the  salt  is  dissolved,  quickly  cool  the  solution,  add  about 
30  cc.  10%  KOH  solution,  and  filter. 

Sodium  cobaltinitrite :  Dissolve  250  g.  NaN02  in  500  cc.  water,  add 
150  cc.  6-normal  HAc  and  25  g.  Co(N03)2.6H20,  let  the 
mixture  stand  over  night,  filter  or  decant  the  solution,  and 
dilute  it  to  one  liter. 

Stannous  chloride,  1-normal:  Dissolve  115  g.  SnCl2.2H2O  in  100  cc. 
12-normal  HC1,  dilute  to  1000  cc.,  and  keep  in  bottles 
containing  granulated  tin. 

Starch  and  potassium  iodide:  Rub  20  g.  starch  to  a  thin  paste  with 
a  little  water  in  a  mortar,  and  pour  the  paste  into  1000  cc. 
boiling  water.  Boil  for  five  minutes,  and  pour  the  liquid 
through  a  funnel  plugged  loosely  with  cotton  wool.  Add  to 
the  filtrate  10  g.  KI  and  5  cc.  chloroform. 

Turmeric:  Shake  turmeric  powder  with  95%  alcohol  and  filter. 

Urea:  Dissolve  200  g.  urea  in  1000  cc.  6-normal  HC1. 


SOLID   REAGENTS. 


Beads  (glass). 

Bismuth    dioxide    (sold    also    as 

sodium  bismuthate). 
Borax  (anhydrous). 
Calcium  chloride  (dry  lumps). 
Copper  (turnings). 
Ferrous  sulfate  (powder). 
Lead  (finely  granulated). 
Potassium  chlorate  (powder). 


Potassium  hydrogen  sulfate. 

Potassium  nitrate. 

Silica  (precipitated). 

Silver  sulfate. 

Sodium  carbonate  (anhydrous). 

Sodium  peroxide  (in  4  oz.  cans). 

Tin  (finely  granulated). 

Zinc  (finely  granulated). 


SOLVENTS. 


Chloroform. 

Ethyl  alcohol  (95%). 

Methyl  alcohol  (acetone-free). 


PREPARATION  OF  THE  TEST-SOLUTIONS. 


Of  the  powdered  salt  whose  formula  is  given  in  the  middle  column  of  the  fol- 
lowing table  weigh  out  the  number  of  grams  given  in  the  last  column,  and  add 
enough  water  (or  acid  when  so  stated  in  the  foot-note)  to  make  the  volume  one 
liter.  To  prepare  the  test-solutions,  dilute  these  stock  solutions,  which  contain 
100  mg.  of  the  constituent  per  cubic  centimeter,  with  nine  times  the  volume  of 
distilled  water.  In  a  few  cases  (indicated  by  the  letter  H)  where  the  substance 
is  not  sufficiently  soluble,  the  stock  solution  is  made  up  so  as  to  contain  50  mg. 
of  the  constituent  per  cubic  centimeter  and  must  be  diluted  with  four  times  its 
volume  of  water  to  yield  the  test-solution. — Since  these  solutions  serve  also 
for  the  preparation  of  the  "unknown  solutions,"  the  purest  salts  that  can  be  pur- 
chased should  be  employed. 


Constit- 

Formula 

Grams        Constit- 

uent 

of  salt 

per  liter 

uent 

Ag 

AgNO3 

160 

Cr 

Pb 

Pb(N03)2 

160 

Zn 

Hg(ous) 

HgNO3.H2O 

140(a) 

Fe(ous) 

Hg(ic) 

Hg(N03)2 

160(a) 

Fe(ic) 

Bi 

Bi(N03)3.5H2O 

230(6) 

Mn 

Cu 

Cu(N03)2.3H2O 

380 

Ni 

Cd 

Cd(N03)2.4H20 

275 

Co 

As(ous) 

As2O3 

13(c) 

Ba 

As(ic) 

As2O5 

150 

Sr 

Sb 

SbCl3 

190(d) 

Ca 

Sn(ous) 

SnCl2.2H2O 

190(e) 

Mg 

Sn(ic) 

SnCl4.3H2O 

270(e) 

Na 

Al 

A1(NO3)3.9H20 

700H 

K 

C03 

Na2C03 

180 

Cl 

80s 

Na2SO3 

160 

Br 

SO4 

K2S04 

90H 

I 

CIO 

NaOCl 

...to) 

SCN 

s 

Na2S.9H20 

750 

N03 

CN 

KCN 

250 

C1O3 

N02 

KNO2 

185 

P04 

Cr04 

K2Cr04 

170 

P04 

Formula 
of  salt 

CrCls(50%sol'n) 

Zn(N03)2 

FeSO4.7H20 

Fe(N03)3.9H2O 

Mn(NO3)2.6H20 

Ni(N03).6H20 

Co(N03)2.6H20 

BaCl2.2H20 

Sr(N03)2 

Ca(N03)2.4H20 

Mg(NO3)2.6H2O 

NaNO3 

KN03 

KC1 

KBr 

KI 

KSCN 

KN03 

KClOa 


Grams 
per  liter 

610 

290 

250H(/) 

715 

530 

500 

500 

180 

240 

590 

530H 

370 

260 

210 
150 
130 
170 
165 
75H 


Na2HP04.12H20  380 
Ca3(P04)2      160(6) 

(a)    Dissolve  in  0.6-normal  HNO«.  (ft)   Dissolve  in  3-normal  HNOi. 

(c)  Digest  with  500  oc.  12-normal  HC1;  then  add  500  cc.  water,  yielding  the  test-solution  of  AsCls 
containing  10  mg.  As  per  cubic  centimeter. 

(d)  Dissolve  in  6-normal  HC1;  and,  in  making  the  test-solution,  dilute  with  2-normal  HC1. 
(«)  Dissolve  in  6-normal  HC1. 

(/)  Dissolve  in  1-normal  H»SO4,  and  keep  in  contact  with  iron  nails. 

(0)   To  600  cc.  of  the  31%  solution  of  commerce  add  450  cc.  water,  yielding  the  test-solution 
containing  10  mg.  CIO  per  cubic  centimeter. 

120 


APPENDICES.  121 

UNKNOWN  SOLUTIONS. 

The  "unknown  solutions"  given  to  the  student  should  contain  the  constituents 
to  be  tested  for  in  quantities  which  are  definitly  known  by  the  instructor.  As 
a  rule  they  may  well  contain  in  10  cc.  300  mg.  of  one  of  the  constituents,  30  mg. 
of  another  of  the  constituents,  and  3  mg.  of  each  of  two  or  three  of  the  remaining 
constituents  of  the  group  in  question.  Such  solutions  may  be  conveniently 
prepared  in  advance  by  mixing  in  a  250  cc.  bottle  60  cc.  of  the  stock  solution 
of  the  first  constituent  (or  120  cc.  if  it  is  half-strength  as  shown  by  an  H  in  the 
table),  6  cc.  of  the  stock  solution  of  the  second  constituent  (or  12  cc.  if  half- 
strength),  and  6  cc.  of  the  test-solutions  of  the  other  constituents,  and  diluting 
with  enough  water  to  make  the  volume  200  cc.  Of  these  "unknown  solutions" 
just  10  cc.  should  be  given  out  to  each  student  for  analysis.  When  time  permits 
two  unknowns  being  done  on  any  group,  the  second  may  well  contain  only  2  mg. 
of  some  of  the  constituents. 


APPARATUS   REQUIRED. 


Returnable. 

4  Lipped  Beakers,  120  cc.  to  500  cc. 
2  Casseroles,  50  cc. 
2  Casseroles,  100  cc. 
2  Casseroles,  200  cc. 
1  Conical  Flask,  30  cc. 
4  Conical  Flasks,  75  cc. 
4  Conical  Flasks,  200  cc.,  with  1  two- 
hole  rubber  stopper. 

1  Conical  Flask,  500  cc. 

2  Jena  Round-Bottom  Flasks,  100  cc. 
with  1  two-hole  rubber  stopper. 

1  Filter  Flask,  500  cc.,  with  a  one-hole 

rubber  stopper. 
1  Flat-Bottom  Flask,  750  cc.,  with  a 

two-hole  rubber  stopper. 
1  Flat-Bottom  Flask,  250  cc.,  with  a 

two-hole  rubber  stopper. 

3  Funnels,  2.5  in. 

1  Funnel,  2  in. 

4  Watch-Glasses,  3  cm. 

2  Watch-Glasses,  7  cm. 
2  Watch-Glasses,  10  cm. 
12  Test-Tubes,  15  cm. 

6  Test-Tubes,  7  cm. 

1  Separating  Funnel,  100  cc. 

1  Graduate,  10  cc. 

1  Graduate,  50  cc. 

1  Nickel  Crucible,  20  cc. 

1  Porcelain  Mortar  and  Pestle. 

1  Desk  Key. 

1  Gas  Burner  (Tin-ill). 

1  Lamp-Stand  and  Rings. 

1  Filter  Stand. 

1  Test-Tube  Rack. 


Not  Returnable. 

12  Hardened  Filters,  5.5  cm. 
1  Pkg.  Filters,  7  cm. 
1  Pkg.  Filters,  9  cm. 
1  Monel-Metal  Wire-Gauze,  12x12  cm. 
1  Box  Labels. 
1  Note-Book. 

1  Piece  Platinum  Foil,  2  cm.  x  1  cm. 
1  Piece  Platinum  Wire,  10  cm. 
1  Horn  Spoon  (bowl  1  cm.  long). 
1  Piece  Glass  Rod,  50  cm.  long. 

1  Sponge. 

2  Towels. 

1  Nichrome  Triangle. 

1  Piece  Rubber  Tubing,  60  cm.  long, 

6  mm.  bore. 
1  Piece  Rubber  Tubing,  60  cm.  long, 

4.5  mm.  bore. 

1  Rubber  Nipple. 

2  Pieces  Glass  Tubing,  each  75  cm. 
long,  4  mm.  bore. 

1  Piece  Hard-Glass  Tubing,  50   cm. 

long,  6  mm.  bore. 
1  Tube  Blue  Litmus  Paper. 
1  Tube  Red  Litmus  Paper. 
1  Box  Matches. 
1  Triangular  File. 
1  Screw  Clamp. 
1  Test-Tube  Brush. 


122 


IONIZATION    VALUES. 


The  following  table  shows  approximately  the  percentage  of  the  substance 
which  is  dissociated  into  its  ions  in  0.1  normal  solution  at  25°.  In  the  case  of 
the  dibasic  acids  the  value  opposit  the  formula  of  the  acid  shows  the  percentage 
of  the  first  hydrogen  that  is  dissociated,  and  that  opposit  the  acid  ion  (HA~) 
shows  the  percentage  of  it  dissociated  (into  H>  and  A-  for  the  case  that  these 
two  ions  are  present  in  equal  quantities). 

Salts  of  type  B+A~  (e.  g.,  KN03) 84% 

Salts  of  type  B+2A=  or  B++A-2  (e.  g.,  K2S04  or  BaCl2) 73 

Salts  of  type  B+3Aa  or  B+++A~3  (e.  g.,  K3Fe(CN)6  or  AlCls) 65 

Salts  of  type  B++A-  (e.  g.,  MgSO4) 40 

KOH,  NaOH 90 

Ba(OH)2 80 

NH4OH 1 

HC1,  HBr,  HI,  HSCN,  HN03,  HC1O3,  HC104,  H2S04,  H2Cr04..  .90 

HsP04,  H3AsO4,  H2S03,  H2C2O4,  HS04~ $0-45 

HN02,  HF 7-9 

HAc,  HCaO,-,  HSO3~ 1-2 

H2S,  H2C03,  H2PO4~,  HOOr 0.1-0.2 

HB02,  HAs02,  HCN,  HC03- 0.002-0.008 

HS-,  HP04- 0.0001-0.0002 

HOH 0.00,000,02 


123 


ATOMIC  WEIGHTS  OF  THE  COMMON  ELEMENTS. 


Aluminum Al 

Antimony Sb 

Arsenic As 

Barium Ba 

Bismuth Bi 

Boron B 

Bromin Br 

Cadmium Cd 

Calcium Ca 

Carbon C 

Chlorin Cl 

Chromium Cr 

Cobalt Co 

Copper Cu 

Fluorin F 

Gold Au 

Hydrogen H 

lodin.  I 


27.1 

120.2 
74.96 

137.37 

208.0 
11.0 
79.92 

112.40 
40.07 
12.00 
35.46 
52.0 
58.97 
63.57 
19.0 

197.2 
1.008 

126.92 


Iron Fe  55.84 

Lead Pb  207.10 

Magnesium Mg  24 . 32 

Manganese Mn  54 . 93 

Mercury Hg  200.6 

Molybdenum Mo  96 . 0 

Nickel Ni  58.68 

Nitrogen N  14.01 

Oxygen O  16.00 

Phosphorus P  31.04 

Potassium K  39.10 

Silicon Si  28.3 

Silver Ag  107.88 

Sodium Na  23.00 

Strontium Sr  87 . 63 

Sulfur S  32.07 

Tin Sn  119.0 

Zinc .  .  Zn  65 . 37 


SOLUBILITIES  OF  SLIGHTLY  SOLUBLE  SUBSTANCES. 


The  numbers  in  the  table  show  the  solubility  in  milli-equivalents  per  liter  at 
20°.  The  letters  v.s.  (very  soluble)  denote  a  greater  solubility  than  1-normal. 
In  the  case  of  the  carbonates  the  values  have  been  corrected  for  hydrolysis  so  as 
to  correspond  to  the  ion-concentration  product  in  the  saturated  solution. — For 
a  general  statement  in  regard  to  the  solubilities  of  other  substances,  see  Note  9, 
page  31. 

Mg         Ca  Sr 

Chloride 
Bromide 
Iodide 
Thiocyanate 
Sulfate 


Chromate 

Carbonate 

Hydroxide 

Fluoride 

Oxalate 

Phosphate 


v.s. 
v.s. 
v.s. 
v.s. 
v.s. 
v.s. 
10. 
0.3 
2.8 
5. 


Ca 

v.s. 

v.s. 

v.s. 

v.s. 

30. 

60. 
0.1 

45. 
0.4 
0.09 
0.7 


v.s. 
v.s. 
v.s. 
v.s. 
1.5 
12. 
0.1 
130. 
1.9 
0.5 


Ba 

v.s. 
v.s. 
v.s. 
v.s. 
0.02 
0.03 
0.1 
450. 
18. 
0.8 


Pb 
70. 
45. 

2.6 
28. 

0.28 

0.0003 

0.0004 

0.2 

5. 

0.012 

0.001 


Ag 

0.01 

0.0005 

0.00002 

0.0008 
50. 

0.16 
*0.2 

0.18 

v.s. 

0.24 

0.05 


124 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 

Return  to  desk  from  which  borrowed. 
This  book  is  DUE  on  the  last  date  stamped  below. 


24Feb'58HK 
REC'D  LD 
1  0  1958 


MAY  22 <5953 

,       nlun'SSRHf 

REC'D  LD 

MAY  2  0-1933 

LD  21-100m-ll,'49(B7146sl6)476 


REG  D 

MAR  2    1962 


SEP 


racut, 


RPR  17 


302162 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


